LABORATORY  COURSE 

IN 

ELECTROCHEMISTRY 


McGraw-Hill  BookCompany 

Purf&s/iers  offioo£§/br 

Electrical  World         The  Engineering  and  Mining  Journal 
Engineering  Record  Engineering  News 

Railway  Age  Gazette  American  Machinist 

Signal  Engineer  American  Engineer 

Electric  Railway  Journal  Coal  Age 

Metallurgical  and  Chem  ical  Engineering  P  owe  r 


LABORATORY  COURSE 


IN 


ELECTROCHEMISTRY 


BY 
OLIVER  P.  ]^ATTS,  PH.  D. 

ASSISTANT   PROFESSOR   OP    APPLIED    ELECTROCHEMISTRY 
THE   UNIVERSITY   OF   WISCONSIN 


FIRST  EDITION 


McGRAW-HILL  BOOK  COMPANY,  INC. 

239  WEST  39TH   STREET,  NEW  YORK 

6  BOUVERIE   STREET,  LONDON,  E.  C. 

1914 


COPYRIGHT,  1914,  BY  THE 
McGRAW-HiLL  BOOK  COMPANY,  INC. 


THE. MAPLE. PRESS. YORK. PA 


PREFACE 

This  laboratory  manual  has  been  designed  primarily 
for  use  in  the  author's  classes  in  the  University  of  Wis- 
consin, and  embodies  the  notes  originally  prepared  by 
C.  F.  Burgess,  former  head  of  the  Chemical  Engineering 
department  of  the  University,  together  with  many  new 
experiments  and  much  additional  material.  It  is  hoped 
that  it  may  prove  a  useful  handbook  in  applied  electro- 
chemical courses  elsewhere  than  at  Wisconsin. 

Thanks  are  due  to  Mr.  C.  F.  Burgess  for  the  use  of  his 
notes,  and  to  Mr.  Claude  N.  Hitchcock  for  the  drawings 
which  illustrate  the  text,  and  for  the  use  of  figures  13, 
14  and  15,  which  were  originally  published  in  his  paper 
upon  Polarization  Single  Potentials  (vol.  25,  Transactions 
of  Americal  Electrochemical  Society) . 

O.  P.  W. 

Sept.,  1914. 


CONTENTS 

PAGE 

INTRODUCTION 1 

LABORATORY  EQUIPMENT 2 

INSTRUCTIONS  FOR  STUDENTS •.  5 

QUALITATIVE  EXPERIMENTS  ON  ELECTROLYSIS     ....  10 

SPECIFIC  RESISTANCE 13 

POLARIZATION 15 

FARADAY'S  LAW 29 

POTENTIAL  AND  ELECTROMOTIVE  FORCE 33 

DISCHARGE  POTENTIALS 56 

OVERVOLTAGE 58 

PASSIVE  STATE 58 

CORROSION  OF  METALS 60 

ELECTROLYTIC  SEPARATION  OF  METALS 63 

ELECTROLYTIC  ANALYSIS 66 

INTERMEDIATE  ELECTRODES 69 

ELECTROPLATING  BATHS 72 

SOLUTIONS  FOR  COLORING  AND  OXIDIZING  METALS.      .      .  79 

PRINCIPLES  OF  ELECTRODEPOSITION 84 

CLEANING  AND  POLISHING 89 

NICKEL  PLATING , 93 

COPPER  PLATING 95 

THE  DEPOSITION  OF  ALLOYS 97 

BRASS  PLATING 100 

SILVER  PLATING 101 

EXPERIMENTS  IN  PLATING 103 

OXIDATION  AND  REDUCTION 121 

OTHER  ELECTROLYTIC  PREPARATIONS 132 

APPENDIX 138 

INDEX  147 


Vll 


A  LABORATORY  COURSE  IN 
ELECTROCHEMISTRY 

INTRODUCTION 

The  experiments  in  this  manual  have  been  chosen  to 
illustrate  the  general  principles  which  underlie  the  more 
important  applications  of  electrochemistry.  No  attempt 
has  been  made  to  adapt  commercial  processes  to  the 
laboratory,  although  a  few  experiments  of  this  character, 
in  which  the  apparatus  lends  itself  readily  to  laboratory 
use,  have  been  included.  The  electric  furnace  and 
batteries  have  been  omitted,  since  they  are  studied  as 
separate  courses  by  the  author's  classes.  Electrolytic 
analysis  has  been  developed  to  such  an  extent  that  it  now 
constitutes  a  study  by  itself,  and  cannot  be  adequately 
treated  in  a  book  on  electrochemistry  in  general.  Al- 
though a  few  simple  experiments  are  given,  the  student 
who  wishes  to  study  this  branch  of  electrochemistry  is 
referred  to  the  well-known  texts  by  Smith  and  Classen. 

Even  with  the  above  limitations  in  scope,  it  will 
probably  be  impossible  for  the  student  to  perform  all 
of  the  experiments  in  the  time  available.  It  will  be 
found,  however,  that  under  many  topics  the  experiments 
are  so  similar  in  nature  that  the  method  of  operation  and 
the  use  of  the  apparatus  may  be  learned  from  any  one  of 
them.  It  is  therefore  suggested  that  in  such  cases  the 
class  or  laboratory  section  work  together  on  the  topic,  a 
particular  experiment  being  assigned  to  one  or  two  stu- 
dents who  report  their  data  and  conclusions  to  the  class 
for  discussion.  All  data  and  results  may  be  recorded  in 

1 


2       A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

the  notebooks,  but  credited  to  the  observer,  so  that  each 
student  will  have  a  record  of  the  experiments  of  others  as 
well  as  of  his  own. 

The  laboratory  work  should  be  accompanied  by  lec- 
tures, recitations  and  assigned  reading  in  books  and  per- 
iodicals. The  latter  is  valuable,  not  only  for  the  infor- 
mation gained,  but  quite  as  much  for  the  acquaintance 
obtained  with  electrochemical  literature  and  the  men  who 
are  producing  it. 

Laboratory  Equipment 

A  brief  description  of  the  laboratory  equipment  which 
is  provided  for  these  experiments  may  be  useful. 

The  laboratory  receives  current  over  three  pairs  of 
cables,  two  from  the  main  switch-board,  and  one  running 
directly  from  the  storage  battery  in  the  basement.  The 
advantage  of  this  direct  connection  of  the  battery  is  that 
experiments  cannot  be  interrupted  by  the  accidental 
pulling  of  the  wrong  plug  from  the  switch-board.  The 
two  pairs  of  cables  from  the  switch-board  are  usually 
connected  to  110  volts  direct  and  alternating  pressure 
respectively.  From  the  cables  wires  are  brought  down 
to  terminals  above  and  at  the  back  of  each  student's 
place  at  the  laboratory  desk.  The  110- volt  direct  cur- 
rent can  be  used  only  during  the  day,  but  the  10-volt 
battery  and  alternating  pressures  are  always  available. 
The  battery  current  is  used  in  the  majority  of  experi- 
ments, while  the  alternating  current  is  useful  for  driving 
motors  and  heating  electrolytes  in  experiments  that  are 
continued  over  night. 

The  storage  battery  consists  of  nine  sets  connected  in 
parallel,  each  set  having  five  160  ampere-hour  cells  in 
series,  making  the  total  capacity  1440  ampere-hours. 
By  means  of  switches,  any  number  of  sets  is  readily 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY       3 

connected  in  series  for  charging.  A  6-volt  battery  of 
half  this  capacity  would  probably  prove  ample  for  ordi- 
nary demands. 

There  is  available  to  each  student  direct  current  at  10 
and  110  volts  pressure.  Control  of  either  current  or 
E.M.F.  as  may  be  required  for  any  particular  experi- 
ment, is  secured  by  portable  rheostats,  consisting  of  small 
lamp  banks,  and  the  ordinary  wire-wound  rheostats  with 
sliding  contacts.  About  a  third  of  the  wire  rheostats 
have  resistances  high  enough  to  permit  their  use  on  the 
110- volt  circuit. 

A  second  pair  of  cables  from  the  battery  would  be  an 
advantage,  as  one  pair  could  then  be  reserved  for  experi- 
ments requiring  constancy  of  line  voltage,  thus  avoiding 
the  fluctuations  produced  by  the  use  of  the  electric 
cleaner  and  similar  intermittent  high-current  work. 

Ammeters,  voltmeters,  rheostats,  etc.,  are  kept  in 
cases  in  the  laboratory,  from  which  they  are  taken  by  the 
students  at  the  beginning  of  the  four-hour  laboratory 
period,  and  returned  at  its  close.  Lamp  cord,  cut  to 
different  lengths,  with  an  inch  of  bare  No.  18  copper  wire 
soldered  to  the  ends,  is  used  for  electrical  connections. 
This  method  of  distribution  and  control  of  current  is 
recommended  for  its  simplicity  and  economy,  and  for 
the  practice  which  it  gives  the  student  in  setting  up 
electrical  circuits  and  in  the  control  of  the  electric 
current.  The  last  feature  is  especially  important. 

Ammeters  having  ranges  of  1  and  of  10  amperes  are 
most  used,  although  at  least  one  instrument  of  50 
amperes  capacity  should  be  available.  Two  or  three 
double-scale  milliammeters  (range  50  and  500)  are 
needed.  Although  double-scale  voltmeters  are  not  so 
convenient  for  rapid  reading  as  single-scale  instruments,  a 
considerable  saving  in  cost  can  be  effected  by  their  use 
for  a  large  part  of  the  voltmeter  equipment.  The  two 


4       A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

scales  should  be  chosen  so  that  both  are  capable  of  direct 
reading,  i.e.,  the  ratio  between  them  should  not  be  3  to  1 
or  30  to  1.  Several  single-scale  instruments  with  a  range 
of  3  volts  are  useful  for  reading  low  voltages  with 
accuracy.  The  medium-priced,  portable  type  of  instru- 
ment is  satisfactory  for  general  laboratory  use,  but 
several  high-grade  instruments  should  be  available  for 
research  work. 

The  outfit  for  measuring  potentials  consists  of  a  two- 
dial  rheostat1  of  1000  ohms  (American  made)  with  its 
coils  connected  for  use  as  a  potentiometer.  This  costs 
$42.50.  A  galvanometer,  keyless  preferred,  costing 
about  $18,  and  a  strong,  well-made,  short-circuit  key 
which  costs  $6.50,  completes  the  outfit.  The  cost  of  this 
outfit,  $67,  is  very  reasonable  when  its  satisfactory  ser- 
vice is  compared  with  that  of  some  cheaper  instruments 
with  which  the  author  has  had  the  misfortune  to  work. 

One  S.P.S.T.,  one  S.P.D.T.,  and  one  D.P.D.T.,  15- 
ampere,  porcelain  base  switch  may  be  fastened  perma- 
nently to  each  desk,  or  they  may  be  mounted  together  on 
a  single  small  board,  which  is  put  away  at  the  end  of 
each  laboratory  period,  leaving  the  desks  clear. 

For  small  electrolytic  cells,  plain,  heavy,  straight  tum- 
blers have  been  found  superior  to  beakers  as  they  are 
rarely  upset  or  broken.  Rectangular  glass  battery  jars 
of  600  and  1400  c.c.  capacity  are  convenient  for  electro- 
plating on  a  small  scale.  For  plating  baths  of  10  to  100 
liters,  rectangular  acid-proof  tanks  may  be  purchased 
very  reasonably  directly  from  the  manufacturers  of  chem- 
ical stoneware. 

For  polishing  metals,  two  or  more  standard  polishing 
lathes  with  cloth  wheels  (bobs)  should  be  provided. 
The  sheet  brass  or  copper  used  for  experiments  in  electro- 
deposition  must  have  a  smooth  surface.  To  avoid  the 

1  Described  on  page  37. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY      5 

trouble  and  expense  of  cutting  down  with  emery,  it  is 
best  to  buy  the  sheet  metal  already  polished. 

An  iron  potash  tank  heated  by  a  steam  coil  is  a  neces- 
sity for  removing  grease  from  large  work.  If  plating  is 
conducted  only  on  a  small  scale,  a  thin  cast-iron  kettle, 
heated  by  a  laboratory  burner,  will  serve  as  the  potash 
tank.  Either  may  be  used  as  an  " electric  cleaner"  by 
connecting  the  tank  directly  to  the  positive  Battery  cable 
by  a  heavy  copper  wire,  and  supporting  across  the  tank, 
but  insulated  from  it,  a  3/8-inch  brass  rod  connected  to 
the  negative  terminal  of  the  battery.  A  pole-reversing 
switch  will  enable  the  object  to  be  made  cathode  or  anode 
at  will. 

Instructions  for  Students 

"  Experience  shows  that  there  is  a  great  tendency  among 
those  who  commence  the  study  of  electrochemistry  to  slop 
through  the  work.  The  average  student  seems  to  think  all 
that  is  necessary  is  to  mix  his  solutions  in  a  more  or  less 
accurate  manner,  and  then  to  switch  on  the  current — the 
electricity  will  do  the  rest.  A  greater  mistake  could  not  be 
made;  unless  details  of  current  density,  electromotive  force, 
temperature  and  composition  of  the  electrolyte  are  carefully 
attended  to,  the  results  will  not  be  such  as  are  expected,  or 
as  are  set  out  in  the  book.  Students  are  very  apt  to  say: 
'It  is  about  right/  or  'The  results  are  near  enough.'  Such 
workers  will  never  succeed  and  do  not  deserve  success." — 
F.  M.  Perkin  in  "Practical  Methods  of  Electrochemistry." 

A  laboratory  experiment  demands  the  best  efforts  of 
the  hand,  eye,  and  brain.  It  is  not  sufficient  to  see  that 
the  apparatus  is  set  up  exactly  as  directed,  that  the  elec- 
trical instruments  are  read  accurately,  that  the  readings 
are  recorded  correctly,  and  that  all  questions  in  the  text 
are  answered  satisfactorily.  Each  experiment,  in  addi- 
tion to  the  purpose  expressed  in  its  title,  should  be  re- 
garded as  an  exercise  to  develop  the  powers  of  observa- 


6       A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

tion,  and  as  a  study  of  the  process  of  electrolysis;  there 
fore  anything  new,  or  of  interest  to  the  student  shoul< 
be  recorded.  To  learn  how  to  set  up  apparatus  and  t 
control  the  electric  current  for  various  purposes  is  also  ; 
valuable  part  of  laboratory  practice. 

Every  electric  circuit  should  contain  a  switch,  in  orde 
that  the  circiiit  may  be  broken  promptly  in  case  of  acci 
dent,  and  a  rheostat  suited  to  the  voltage  employed 
Knowing  the  line  voltage  and  the  current  desired,  th 

resistance  required  for  the  rheostat  may  be  calculate< 

•p 

from  the  equation  representing  Ohm's  Law:  R  =  y   1] 

which  E  is  the  E.M.F.  in  volts,  I  is  the  current  ii 
amperes,  and  R  is  the  resistance  in  ohms.  In  connect 
ing  a  rheostat,  great  care  should  be  taken  to  see  that  a 
the  outset  it  is  adjusted  for  its  maximum  resistance 
Carelessness  in  this  respect  may  result  in  the  injury 
or  even  in  the  destruction  of  the  electrical  instruments 

Students  are  expected  to  report  at  once  any  damagi 
to  instruments,  and  are  held  responsible  for  repairs. 

The  polarity  of  all  cable  terminals  in  the  laboratory 
should  be  plainly  marked,  and  at  least  one  terminal  o 
every  ammeter  and  voltmeter  has  been  marked  by  th< 
maker.  Ammeters  should  be  connected  in  series,  i.e. 
so  that  the  full  current  flowing  in  the  circuit  must  pas; 
through  the  instrument.  Unless  otherwise  specifically 
stated,  voltmeters  should  be  connected  across  the  twc 
points  whose  difference  of  potential  it  is  desired  to  meas 
ure.  The  positive  terminal  of  the  ammeter  should  b< 
connected  to  the  wire  which  has  come  from  the  positive 
terminal  of  the  line,  and  the  positive  of  the  voltmetei 
should  go  to  the  anode,  i.e.,  to  the  side  of  the  electrolytic 
cell  which  is  connected  to  the  positive  of  the  line. 

Before  closing  the  switch,  the  student  should  trace  ou1 
the  circuit  carefully,  to  see  that  the  current  will  go  onlj 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY      7 

where  desired,  and  that  the  instruments  are  connected  so 
that  the  current  will  enter  at  their  positive  terminals. 
This  is  best  done  by  starting  at  the  positive  terminal  of 
the  line  and  following  the  circuit  by  actual  physical  con- 
tact of  the  hand,  through  the  cell  and  instruments, 
around  to  the  negative  terminal  of  the  line.  For  taking 
down  the  apparatus  there  is  but  one  safe  rule  to  follow: 
after  opening  the  switch,  first  disconnect  both-  wires  from 
the  line  terminals.  This  insures  that  no  live  wires  are  on 
the  desk  top,  and  all  danger  of  a  short-circuit  is  avoided. 

In  order  to  use  electrical  instruments  intelligently  and 
safely,  it  is  necessary  to  understand  their  construction. 
The  principles  of  operation  of  the  Weston  direct-current 
instruments  will  be  described. 

The  millivoltmeter  is  the  basis  of  both  the  voltmeter 
and  the  ammeter.  Examination  shows  that  the  milli- 
voltmeter consists  of  a  movable  coil  of  very  fine  wire,  to 
which  the  needle  is  attached,  carefully  pivoted  on  jew- 
elled bearings  between  the  poles  of  a  permanent  magnet. 
Current  enters  and  leaves  the  coil  through  very  fine  steel 
springs,  one  of  which  is  visible  at  the  upper  bearing. 
When  current  passes  through  the  coil,  it  sets  up  a  mag- 
netic field,  which,  reacting  with  the  field  of  the  permanent 
magnet,  turns  the  coil  and  moves  the  needle  up  the  scale. 
When  the  current  ceases,  the  springs  return  the  needle  to 
zero.  The  delicacy  of  the  springs  and  coil  indicate  that 
only  extremely  small  currents  can  be  sent  through  a  milli- 
voltmeter without  injuring  it. 

The  voltmeter  consists  of  a  millivoltmeter  with  a 
resistance  placed  in  series  with  the  movable  coil.  The 
scale  is  marked  in  volts,  and  the  resistance  is  made  such 
that  full-scale  deflection  is  obtained  when  the  voltage 
for  which  the  instrument  is  designed  is  impressed  across 
its  terminals.  Compare  the  resistance  of  a  millivolt- 
meter  with  the  resistances  of  several  voltmeters  of  differ- 

2 


8       A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

ent  ranges.  Compute  the  current  which  flows  through 
each  instrument  at  full-scale  deflection.  In  the  double 
scale  voltmeter,  the  series  resistance  is  divided  into  two 
portions,  and  a  third  binding  post  is  placed  between  the 
resistance  coils.  What  must  be  the  relation  between  the 
total  resistances  of  the  instrument  when  using  the  upper 
and  the  lower  scales,  if  one  scale  is  twice  the  other? 
Fifty  times?  Verify  your  prediction  by  looking  up  the 
resistances  marked  on  laboratory  instruments. 

The  ammeter  is  made  up  of  a  millivoltmeter,  and  a 
carefully  made  resistance  called  a  shunt.  The  current 
to  be  measured  is  sent  through  the  shunt,  and  the  milli- 
voltmeter is  connected  so  that  it  reads  the  fall  of  potential 
produced  in  the  shunt  by  the  current.  The  scale  is 
usually  marked  directly  in  amperes.  For  currents  less 
than  100  amperes,  the  shunt  is  often  concealed  within  the 
case  of  the  instrument,  but  for  very  large  currents,  it  is 
always  separate  from  the  millivoltmeter,  and  is  then 
called  an  external  shunt.  It  is  evident  that,  by  the  use  of 
a  second  shunt  having  a  different  resistance,  the  range  of 
the  instrument  may  be  greatly  altered. 

Danger!  The  millivoltmeter,  usually  marked  amme- 
ter, designed  for  use  with  an  external  shunt,  must  never 
be  directly  connected  to  the  electrical  circuit. 

Voltmeters  may  be  protected  from  injury  by  taking 
care  to  connect  them  to  the  circuit  so  that  the  deflection 
of  the  needle  is  in  the  proper  direction,  and  by  never 
connecting  them  to  a  source  of  pressure  greater  than  the 
scale  reading.  When  in  doubt  whether  a  circuit  is  of 
high  or  low  pressure,  use  a  high  reading  instrument. 

Ammeters  should  be  protected  by  a  rheostat  of  such 
resistance  that  the  current  will  be  within  the  range  of  the 
ammeter.  Fuses  may  of  course  be  used,  but  they  are 
not  necessary  for  careful  students. 

The  number  and  range  of  voltmeters  and  ammeters 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY       9 

used  should  be  recorded.  The  use  of  instruments  of 
ranges  unsuited  to  the  experiment  may  give  inaccurate 
results,  or  one  of  the  instruments  may  be  at  fault.  In 
the  latter  case,  a  calibration  of  the  instrument,  and  the 
substitution  of  the  correct  values  will  avoid  the  necessity 
of  repeating  the  experiment. 

The  notebook  should  be  about  7X8  1/2  inches,  ruled 
in  small  squares.  Experiments  should  be  nnmbered  and 
dated,  and  each  should  begin  at  the  top  of  a  page,  with 
the  title,  or  a  statement  of  the  purpose  of  the  experiment. 
The  original  data,  and  the  calculated  results  should  be 
recorded  in  a  single  table.  The  equation  by  which  the 
results  are  calculated  should  be  given,  together  with  a 
sample  substitution  of  numerical  values  in  the  equation. 
This  will  assist  the  instructor  in  locating  mistakes  in  cal- 
culation. Observations  made  during  the  experiment 
may  be  made  below  the  table,  or  on  the  opposite  page, 
but  an  endeavor  should  be  made  to  confine  the  experi- 
ment to  two  pages  that  face. 

Current  densities  are  given  in  amperes  per  square  deci- 
meter unless  otherwise  stated. 

Laboratory  Experiments 

The  passage  of  a  unidirectional  current  through  an 
electrolyte  is  accompanied  by  the  liberation  of  some  sub- 
stance at  each  electrode.  If  the  liberated  substances  are 
capable  of  combining  with  the  electrodes  under  the  con- 
ditions of  the  experiment,  they  do  so.  If  incapable  of 
combining  with  either  the  electrodes  or  the  electrolyte, 
they  are  set  free  at  the  electrodes . 

The  first  seven  experiments  are  qualitative  in  nature 
and  are  intended  to  give  familiarity  with  the  general 
nature  of  the  process  of  electrolysis.  Each  student 
should  perform  them  all.  Their  satisfactory  perfor- 


10     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

mance  requires  constant  attention  and  keen  use  of  the 
eyes  to  observe  all  that  happens.  Phenomena  demand- 
ing attention  are  the  solubility  or  insolubility  of  anodes, 
the  evolution  of  gas  at  either  or  both  electrodes,  what 
classes  of  substances  are  liberated  at  the  anode  and  what 
at  the  cathode,  and  the  effect  of  low-  and  of  high-current 
densities  (0.2  to  5  amperes  per  sq.  dm.)  The  questions 
with  each  experiment  are  intended  as  suggestions,  not  as 
limitations  to  observation  by  the  student.  Litmus  paper 
will  prove  useful  in  detecting  changes  which  might  other- 
wise be  overlooked.  The  reactions  occurring  at  the  elec- 
trodes or  in  the  electrolyte  may  be  expressed  most  briefly 
by  chemical  equations. 

EXPERIMENT  1 

THF    EFFECT    OF    DIFFERENT    ELECTRODES    UPON    THE 
PROCESS  OF  ELECTROLYSIS 

Using  about  an  eight  percent  solution  of  copper  sul- 
phate, connect  three  cells  in  series. 

ANODE  CATHODE 

a.  copper  carbon 

b.  carbon  carbon 

c.  lead  lead 

Pass  a  current  of  about  0.5  amperes  per  sq.  dm.  and 
observe  the  effect  upon  the  electrodes  and  the  electrolyte. 
Increase  the  current  to  three  or  four  times  its  initial 
value,  and  after  five  minutes  again  observe  the  effects. 
What  electrodes  are  insoluble  in  this  electrolyte?  What 
is  the  gas? 

Similarly  use  dilute  sulphuric  acid  as  electrolyte. 

ANODE  CATHODE 

d.  lead  lead 

e.  copper  lead 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     11 

EXPERIMENT  2 

REACTIONS  BETWEEN  ELECTRODE  PRODUCTS 
Pass  current  through  three  cells  in  series. 

ANODE  CATHODE  ELECTROLYTE 

a.  lead           lead  NaN03  15  percent 

6.  iron            iron  NaNOs  15  percent 

c.  iron            iron  NaCl      15  percent 

What  anodes  are  insoluble?  Explain  the  formation  of 
precipitates.  Why  the  different  results  in  b  and  c? 
What  are  the  gases?  How  do  you  explain  the  surprising 
phenomenon  which  occurs  in  6  at  low-current  densities? 
Can  current  pass  without  equivalent  chemical  change? 

EXPERIMENT  3 
ELECTROLYTIC  REDUCTION 

Use  a  lead  anode,  and  a  cathode  of  galena  (PbS)  in 
dilute  sulphuric  acid.  What  chemical  changes  occur? 

EXPERIMENT  4 
ELECTROLYSIS  OF  POTASSIUM  BROMIDE 

Use  an  anode  of  platinum  wire,  and  a  cathode  of  copper 
in  a  ten  percent  solution  of  potassium  bromide.  Place 
the  electrodes  far  apart,  and  test  with  litmus  paper 
between  as  well  as  at  the  electrodes.  Use  a  very  small 
current  at  first.  Increase  this  until  there  is  a  change  in 
the  action  occurring  at  the  anode.  Such  liberation  of 
gas  at  anode  or  cathode  may  usually  be  brought  about  by 
the  use  of  extremely  high-current  densities.  How  do  you 
explain  it? 


12     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 
EXPERIMENT  5 

THE  EFFECT  OF  THE  CATHODE  MATERIAL  UPON 
ELECTROLYSIS 

Electrolyze  a  ten  percent  solution  of  salt,  using  a  very 
small  (0.05  ampere)  current  for  the  first  three  minutes, 
and  the  two  cells  in  series. 

a.  With  carbon  electrodes. 

6.  With  a  carbon  anode  and  a  mercury  cathode. 
Pour  a  thin  layer  of  mercury  into  the  electrolytic  cell  and 
make  contact  with  it  by  means  of  an  insulated  wire. 

What  difference  is  noted  in  the  two  cells  on  starting 
electrolysis?  Is  it  a  case  of  the  passage  of  current  with- 
out chemical  change  at  one  electrode?  After  five  min- 
utes, wash  the  mercury  free  from  the  salt  solution,  place 
5  c.c.  distilled  water  on  it  and  set  aside  in  contact  with 
litmus  paper.  Repeat  the  experiment  with  a  fresh  por- 
tion of  mercury  using  a  current  of  about  1  ampere  to 
hasten  action. 

The  mercury  cathode  is  extensively  used  in  commercial 
cells  for  the  electrolysis  of  sodium  chloride. 

EXPERIMENT  6 
THE  ELECTRO-DEPOSITION  OF  LEAD 

With  lead  electrodes  electrolyze  a  fifteen  percent  solu- 
tion of  lead  nitrate  or  lead  acetate.  Dilute  the  solution 
with  three  volumes  of  water  and  repeat. 

EXPERIMENT  7 

AN  EXAMPLE  OF  THE  MAKING  OF  CHEMICAL  COMPOUNDS 
ELECTROLYTICALLY 

In  Luckow's  patented  process  for  the  manufacture  of 
white  lead,  the  electrolyte  consits  of  13  g.  sodium  chlorate 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     13 

and  2  g.  sodium  carbonate  per  liter.  Carbon  dioxide 
diluted  by  air  is  passed  in  at  the  cathode. 

Use  lead  electrodes  with  0.25  amperes  per  sq.  dm.  in 
this  electrolyte,  stirring  occasionally  by  hand,  or  con- 
tinuously by  a  jet  of  air,  for  a  quarter  of  an  hour.  The 
carbon  dioxide  may  be  omitted  in  this  brief  experiment. 
Examine  the  product  for  pigment  purposes. 

Now  repeat  the  experiment  using  as  electrolyte  a  so- 
lution of  15  g.  sodium  carbonate  per  liter.  Look  up  the 
solubilities  of  lead  carbonate,  lead  chlorate  and  lead 
hydroxide,  if  you  do  not  already  know  them. 

What  is  the  use  of  the  sodium  chlorate?  Since  a  basic 
carbonate  of  lead  is  the  product  desired,  and  sodium  car- 
bonate is  necessary  for  its  continuous  production,  is  it  not 
strange  that  so  little  sodium  carbonate  is  used  in  the 
electrolyte.  Explain. 

The  claim  is  made  for  this  process  that  the  electrolyte 
is-not  used  up,  that  white  lead  can  be  produced  continu- 
ously by  hanging  in  lead  anodes,  blowing  in  carbon  diox- 
ide and  passing  the  current.  Can  you  demonstrate  this 
claim  by  equations  of  the  various  reactions? 

Classify  as  soluble  or  insoluble  the  anodes  so  far  used. 

Specific  Resistance  or  Resistivity  of  Electrolytes 

In  commercial  electrolysis  it  is  desirable  that  the 
resistance  of  the  electrolytes  used  be  as  low  as  possible,  in 
order  to  minimize  the  transformation  of  electrical  energy 
into  heat  by  the  passage  of  the  current  through  the 
electrolyte.  The  energy  so  transformed  is  directly  pro- 
portional to  the  resistance  of  the  electrolyte,  and  to  the 
square  of  the  current,  and  is  usually  referred  to  as  the 
PR  loss.  One  of  the  first  questions  to  be  answered  con- 
cerning any  electrolyte  proposed  for  commercial  elec- 
trolysis is,  ''What  is  its  resistivity?"  Resistivities  are 


14  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 


expressed  in  ohms  per  centimeter  cube,  and  may  be 
conveniently  determined  by  readings  of  current  and 
fall  of  potential  across  the  electrolyte  contained  in  a 
vessel  of  uniform  and  known  cross  section  and  length. 
It  is  obviously  necessary  that  the  electrodes  completely 
fill  the  ends  of  the  vessel.  Since  the  resistance  of  elec- 
trolytic conductors  is  greatly  affected  by  temperature, 
the  thermometer  should  be  read  for  each  determination. 
To  obtain  the  total  resistance  of  the  solution,  apply 

Ohm's  law — R  =  — •     The  resistivity    is    obtained    by 
dividing  the    total  resistance  by  the  .length  in  centi- 


Battery 


Cell 


R 
V\AAA/\AAA/ 


FIG.  1. 

meters,  and  multiplying  by  the  area  in  square  centi- 
meters. Why?  A  simple  way  of  obtaining  the  cross 
section  is  to  measure  the  volume  of  electrolyte,  and  divide 
this  by  its  length.  Polarization  may  cause  serious  errors 
unless  special  precautions  are  taken. 

In  the  experiments  immediately  following,  it  is  desir- 
able for  comparison  of  the  resistivities  of  different  elec- 
trolytes and  of  the  same  electrolyte  measured  by  different 
methods,  that  readings  be  made  at  some  fixed  tempera- 


A  LABORATOEY  COURSE  IN  ELECTROCHEMISTRY     15 

ture  in  all  experiments,  25°  C.  is  suggested.  It  may  be 
of  interest  to  obtain  readings  at  other  temperatures  also. 
The  continued  passage  of  current  will  raise  the  tempera- 
ture of  the  electrolyte. 

The  voltmeter,  ammeter,  rheostat,  switch  and  electro- 
lytic cell  may  be  connected  as  shown  in  Fig.  1. 

EXPERIMENT  8 

THE  RESISTIVITY  OF  A  NORMAL  SOLUTION  OF  COPPER 
SULPHATE 

In  a  rectangular  vessel  about  10  cm.  long,  and  15  to 
20  sq.  cm.  in  section,  put  120  to  150  c.c.  of  solution. 
Use  electrodes  of  clean  sheet  copper,  an  ammeter  of  about 
1  ampere  capacity,  a  rheostat  and  voltmeter  suited  to 
whatever  low  pressure,  5  to  10  volts,  may  be  available. 
Connect  as  in  Fig.  1.  Measure  the  distance  between 
electrodes.  For  this  experiment  secure  one  reading  at  a 
current  below  0.1  ampere,  and  one  above  0.5  ampere  at 
the  same  temperature.  Having  secured  all  data  desired, 
replace  the  copper  electrodes  by  clean  sheets  of  lead  and 
repeat. 

Does  the  voltmeter  needle  return  to  zero  at  once  in  each 
case  on  opening  the  switch?  If  not,  record  the  highest 
reading  each  time.  This  is  the  polarization  referred  to 
later. 

Compute  the  total  resistance  and  the  resistivity. 

Is  there  any  choice  as  regards  accuracy  between 
the  use  of  a  current  of  0.1  and  0.5  amperes  in  this 
experiment? 

Polarization 

The  E.M.F.  observed  in  some  cases  when  the  switch 
is  opened  is  due  to  polarization.  Since  there  is  a  complete 
electrical  circuit  through  the  voltmeter,  this  E.M.F.  must 


16     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

send  a  current  through  the  electrolytic  cell.  What  is  the 
direction  of  this  current  compared  to  that  of  the  original 
current  when  the  switch  was  closed?  It  is  for  this  reason 
that  the  E.M.F.  which  results  from  the  passage  of  current 
through  an  electrolytic  cell  is  often  called  the  Counter 
Electromotive  Force  of  Polarization.  Its  cause  is  the  same 
as  that  of  the  E.M.F.  of  any  voltaic  cell,  which  may  con- 
sist of  two  unlike  electrodes  making  contact  with  a  single 
electrolyte,  or  even  of  like  electrodes  in  two  different 
electrolytes,  i.e.,  an  unsymmetrical  electrochemical 
system.  Initially,  many  electrolytic  cells  consist  of 
like  electrodes  in  a  uniform  electrolyte,  as  in  the  last 
experiment.  If  the  passage  of  current  causes  polariza- 
tion, investigation  will  show  that  the  system  has  become 
unsymmetrical  by  a  change  in  at  least  one  electrode,  or  a 
change  in  the  material,  concentration,  or  temperature  of 
the  film  of  electrolyte  in  contact  with  the  electrode. 

In  all  future  experiments  where  polarization  is  noted, 
the  student  should  observe  whether  it  is  due  to  a  change  in 
the  material  of  the  electrode,  or  electrolyte,  or  to  a  con- 
centration change  in  the  latter.  This  is  important  for 
the  attainment  of  a  practical  knowledge  of  electrochemi- 
cal phenomena. 

Correction  for  Errors  Due  to  Polarization 

In  calculating  the  resistance  of  electrolytes  by  Ohm's 
law,  it  was  assumed  that  the  E.M.F.  observed  was  entirely 
spent  in  forcing  current  through  the  electrolyte,  but  it  is 
evident  that  part  of  it  was  offset  by  the  counter  E.M.F. 
of  polarization,  so  that  the  equation  should  be: 

*=¥• 

Recalculate  the  former  values  of  resistivity,  correcting 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     17 

for  polarization.     What  was  the  percent  of  error  in  the 
unconnected  values? 

EXPERIMENT  9 

THE  RESISTIVITY  OF  NORMAL  COPPER  SULPHATE  USING 
A  HIGH  E.M.F. 

Using  a  cell  30  to  40  cm.  long,  and  8  sq.  cm.  in 
cross  section,  measure  the  resistivity  of  normal  copper 
sulphate  and  normal  sulphuric  acid  solutions  with 
lead  electrodes  and  a  110- volt  source  of  pressure. 
Take  care  that  a  suitable  rheostat  is  employed 
and  that  it  is  adjusted  for  its  highest  resistance  before 
closing  the  switch.  An  application  of  Ohm's  law  will 
tell  whether  the  rheostat  selected  is  sufficient  to  reduce 
the  current  within  the  range  of  the  1  ampere  ammeter 
used.  What  is  the  percent  of  error  in  resistivity  due  to 
polarization? 


-•  no  v- 


FIG.  2. 

If  it  is  desired  to  read  the  polarization  with  greater 
accuracy  than  can  be  done  with  the  high-range  voltmeter 
used,  the  diagram  in  Fig.  2  may  be  employed  for  con- 
necting another  low-reading  voltmeter  by  means  of  a 
double-throw  switch  so  that  this  is  not  in  the  circuit 
until  after  the  line  switch  has  been  opened.  A  double- 
scale  voltmeter  may  be  used  instead  of  two  instruments 


18     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

by  connecting  as  in  Fig.  3.  If  the  instrument  has  two 
positive  terminals,  as  is  the  case  with  some  makes,  they 
must  be  connected  to  the  switch  and  to  the  anode, 
instead  of  to  the  cathode  as  shown.  This  involves  trans- 
ferring the  switch  to  the  anode  side  of  the  cell. 


s, 

72 

u", 

V 

3  t 

A 

/t 

Pol 

1 

+    ' 

4 

FIG.  3. 


EXPERIMENT  10 

THE  RESISTIVITY  OF  METALS  AND  ALLOYS 

Measure  the  resistivity  of  3  or  4  feet  of  fine  wire. 
Between  Nos.  22  and  30,  copper,  iron,  monel  and 
nichrome  may  be  used.  The  current  should  be  small 
enough  to  cause  no  sensible  heating.  A  millivoltmeter 
may  be  substituted  for  the  voltmeter  in  any  case  where 
the  latter  has  proved  of  too  high  a  range.  A  low-voltage 
circuit  should  be  used,  and  since  the  ordinary  rheostats 
are  not  of  sufficient  resistance  for  use  in  series,  they  may 
be  used  in  shunt,  as  shown  in  Fig.  4. 

The  fullline  voltage  is  impressed  across  the  rheostat, 
then  by  the  sliding  contacts  any  desired  fraction  of  this  is 
picked  off  for  use.  This  method  is  useful  whenever  it  is 
desired  to  increase  the  impressed  E.M.F.  from  zero  up 
by  small  increments. 

Measure  the  diameter  of  the  wire  by  a  micrometer  and 
compute  the  resistivity.  Determine  whether  the  resis- 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     19 

tivity  increases  or  diminishes  with  rise  of  temperature, 
by  passing  a  current  large  enough  to  heat  the  wire  appre- 
ciably; compare  the  result  with  the  resistivity  at  the 
lower  temperature.  A  conductor  whose  resistance 
increases  with  rise  of  temperature  has  a  positive  tempera- 
ture coefficient  of  resistance.  Do  you  find  that  of 
metals  and  alloys  to  be  positive  or  negative?  That  of 
electrolytic  conductors? 


EXPERIMENT  11 
RESISTIVITY  OF  ELECTROPLATING  SOLUTIONS 

Measure  the  resistivity  of  several  of  the  following 
plating  solutions,  either  taken  from  the  laboratory  plating 
tanks,  or  made  up  according  to  the  formulas  on  pages 
72-79. 

1.  Nickel  solution. 

2.  Brass  solution — a  deadly  poison! 

3.  Copper  cyanide  solution — a  deadly  poison! 

4.  Acid  copper  sulphate  solution. 

5.  Zinc  solution. 


20     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

Avoid  polarization  by  the  use  of  a  soluble  anode  like  the 
metal  deposited,  or  ascertain  its  amount  (pages  16-18), 
and  correct  for  it.  If  soluble  electrodes  are  used,  they 
must  be  the  same  as  the  metal  in  solution  in  order  that 
the  bath  may  not  be  spoiled  by  the  introduction  of  a 
foreign  metal. 

Correction  for  the  Voltmeter  Current 

A  reference  to  Fig.  1,  page  14,  will  show  that  the 
ammeter  reads  the  current  flowing  through  the  voltmeter 
in  addition  to  that  passing  through  the  electrolytic  cell. 
Look  up  the  resistance  of  the  voltmeter*  used  in  experi- 
ment 11,  and  by  applying  Ohm's  law,  compute  the  volt- 
meter current  for  several  of  your  readings.  To  what 
extent  does  this  error  affect  the  resistivities  previously 
computed?  Is  it  desirable  to  apply  the  correction  or 
not?  Why? 

This  error  in  the  data  may  be  avoided  by  connecting 
the  voltmeter  so  that  it  includes  both  the  cell  and  the 
ammeter.  Draw  a  diagram  showing  this  connection. 
This  arrangement  should  be  used  in  measurements 
involving  currents  of  a  few  hundredths  of  an  ampere,  or 
else  the  correction  for  the  voltmeter  current  should  be 
applied  to  the  current  readings. 

EXPERIMENT  12 

MEASUREMENT  OF  THE  RESISTANCE  OF  A  VOLT- 
METER 

The  resistance  of  a  voltmeter  is  usually  marked  upon 
the  instrument  by  the  maker,  but  this  may  become  oblit- 
erated, so  that  it  is  sometimes  necessary  to  measure  the 
resistance  of  a  particular  voltmeter.  If  another  volt- 
meter of  about  the  same  range  and  known  resistance  is 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     21 

available,  the  two  are  connected  in  series  in  a  suitable 
circuit,  and  the  readings  recorded.  Since  the  fall  of 
potential  over  different  parts  of  any  electrical  circuit  is 
proportional  to  the  resistances,  it  is  evident  that  the 
resistances  of  the  two  instruments  are  in  the  same  ratio 
as  their  respective  readings. 

Measure  the  resistances  of  two  voltmeters  against  that 
of  a  standard  instrument. 


EXPERIMENT  13 

THE  MEASUREMENT  OF  RESISTANCES  BY  USE  OF  A  VOLT- 
METER ONLY 

Resistances  may  be  measured  without  the  use  of  an 
ammeter  by  the  connections  shown  in  Fig.  5. 

The  voltmeter  used  must,  of  course,  be  adapted  to  the 
line  pressure  employed.  First  throw  the  S.P.D.T.  switch 


Line    + 


E.M.F. 


Cell 


FIG.  5. 

"S"  to  the  left,  and  read  the  line  voltage  E,  then  throw 
the  switch  to  the  right,  connecting  the  cell  in  series  with 
the  voltmeter,  and  read  E'.  The  resistance  of  the  cell  is 

E  —  E' 

found  from  the  formula  R  =  R'  X  — — — >   in  which 

hi 

R'  is  the  resistance  of  the  voltmeter.  The  principle  is 
that  of  experiment  12.  E  —  E'  is  the  fall  of  potential 
over  what? 


22     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

For  respectable  accuracy  it  is  necessary  that  E'  lie 

E          2E 

between  ^  and  -^-  in  value.  The  length  and  cross  sec- 
tion of  the  electrolyte  may  be  varied  to  secure  this  result. 
It  may  even  be  necessary  to  change  to  a  voltmeter  of 
higher  or  lower  resistance  (probably  necessitating  a 
change  in  line  voltage)  to  secure  the  above  condition. 
The  use  of  a  high  E.M.F.  lessens  the  error  caused  by 
polarization.  Why? 

Study  this  method  by  measuring  the  resistivity  of 

a.  Normal  copper  sulphate. 

6.  Vioo  normal  copper  sulphate. 

c.  Hydrant  water. 

EXPERIMENT  14 

MEASUREMENT  OF  RESISTANCE  BY  USE  OF  VOLTMETER 
AND  RESISTANCE  Box 

If  a  resistance  box  is  available,  the  principle  of  ex- 
periment 13  may  be  more  conveniently  applied  by 
connecting  as  in  Fig.  6. 


Line      -4- 


FIG.  6. 


S'  is  a  D.P.D.T.  switch  for  connecting  the  voltmeter 
successively  across  the  resistance  box  and  the  cell.  The 
resistance  should  be  adjusted  until  the  voltmeter  reading 
across  each  is  the  same;  then  the  resistance  of  the  cell 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     23 


equals  that  of  the  box,  except  for  the  error  due  to  polariza- 
tion. By  opening  S  when  S'  is  closed  to  the  left,  the 
polarization  may  be  read.  Having  determined  the 
polarization,  it  is  easy  to  correct  for  it.  How? 

By  this  method  measure  the  resist;vity  of  a  nickel 
plating  bath  and  of  hydrant  water.  Add  one  part  of  the 
plating  solution  to  twenty-four  parts  of  water,  and  re- 
measure.  What  error  is  caused  in  each  by  "polarization 
when  a  correction  is  not  applied  for  it? 

EXPERIMENT  15 

THE    TEMPERATURE   COEFFICIENT   OF   RESISTANCE    OF 
AN  ELECTROLYTE 

By  means  of  the  glass  cell  shown  in  Fig  7,  measure 
the  resistance  of  normal  copper  sulphate  between  room 


FIG.  7. 

temperature  and  75°  C.  taking  readings  at  10°  intervals. 
Current  is  sent  through  the  cell  by  means  of  copper 
electrodes,  and  the  fall  of  potential  over  a  constant  length 
of  electrolyte  is  read  on  the  voltmeter  attached  to  small 
copper  wires  introduced  through  the  tubes  A  and  B. 
Temperatures  are  read  on  the  thermometer  T.  A  1- 
ampere  meter  and  suitable  rheostat  should  be  connected 
in  series.  Using  a  110- volt  circuit  this  particular  elec- 


24     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

trolyte  will  be  heated  sufficiently  by  the  current.  To 
avoid  a  rise  of  temperature  while  the  readings  are  being 
made,  it  is  suggested  that  0.6  to  0.7  ampere  be  used  for 
heating,  and  that  this  be  reduced  to  0.2  ampere  when  it 
is  desired  to  take  a  reading.  This  change  may  be  made 
conveniently  by  the  use  of  a  double-throw  switch.  The 
small  lamp  rheostats  are  already  provided  with  this. 

If  air  bubbles  are  caught  in  the  tube  when  filling  it, 
what  will  be  their  effect  upon  the  results?  Why?  Watch 
for  the  collection  of  air  in  the  tube  at  the  higher  tempera- 
tures also.  Plot  the  results  in  the  form  of  a  curve,  and 
determine  the  temperature  coefficient,  a,  for  the  inter- 
vals 30°  to  40°  and  60°  to  70°,  from  t'he  formula  R'  = 
R°  (1  +  at). 

F.  Kohlrausch's  method  of  measuring  the  resistance 
of  electrolytes  by  means  of  alternating  current,  using 
the  slide- wire  bridge  and  a  telephone  receiver,  completely 
eliminates  polarization,  but  has  special  errors  of  its  own, 
so  that  on  the  average  it  is  but  slightly  more  accurate 
than  the  fall-of-potential  method.  This  method  is 
described  in  Ostwald-L/uther's  Physiko-Chemische  Mes- 
sungen,  2nd  edition,  pages  395-411. 

EXPERIMENT  16 
THE  EFFECT  OF  THE  SOLVENT  UPON  RESISTIVITY 

Previous  experiments  have  shown  that  the  resistivity 
of  electrolytes  varies  with  the  nature  of  the  dissolved 
substance,  and  with  its  amount.  The  solvent  also  has  a 
marked  effect  upon  resistivity. 

Dissolve  15  g.  copper  nitrate  in  150  c.c.  of  distilled 
water,  also  in  the  same  amount  of  denatured  alcohol. 
By  one  of  the  previous  methods  measure  the  resistivity 
of  each  solution. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     25 


EXPERIMENT  17 

RESISTANCE  OF  A  FUSED  ELECTROLYTE 

Support  a  small  iron  crucible  with  a  binding-post 
attached  to  it  over  a  Bunsen  burner.  Make  a  flat 
spiral  of  No.  18  B.  &  S.  gauge  iron  or  nichrome  wire, 
bend  the  end  in  the  center  up  at  right  angles  to  the  spiral 
and  suspend  this  as  one  electrode  a  half  inch  above  the 
bottom  of  the  crucible.  Rigidly  support  a  thermo 
couple  inside  the  crucible,  and  connect  it  to  its  milli- 
voltmeter.  Connect  a  double  scale  0-15-150  voltmeter 


no  v* 


• WWWWVX 


FIG.  8. 


across  the  electrodes,  and  a  low-reading  ammeter  and  a 
lamp  bank  in  series  with  the  110-volt  circuit  as  indicated 
in  Fig.  8. 

The  switch  S'  which  connects  the  low  scale  of  the  volt- 
meter to  the  circuit  should  be  closed  only  at  the  instant 
of  taking  a  reading,  and  then  only  when  the  reading  on 
the  high  scale  shows  the  voltage  to  be  well  within  the 
range  of  the  low  scale.  Put  some  pulverized  sodium 
nitrate  in  the  crucible,  light  the  burner,  and  add  more 


26     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

sodium  nitrate  as  it  melts.  When  enough  is  melted,  be- 
gin taking  readings  at  one  minute  intervals.  Record 
time,  amperes,  volts,  and  millivolts.  When  heated 
sufficiently,  shut  off  the  gas  and  take  readings  as  it  cools. 
Solidification  should  be  indicated  by  a  slight  lag  in  the 
rate  of  cooling. 

Plot  the  temperature-resistance  curve,  and  note  the 
location  of  the  point  of  sharpest  inflection  as  compared 
with  the  freezing  temperature. 

Sodium  hydroxide  may  be  used  in  place  of  the  nitrate. 

EXPERIMENT  18 

i 

THE  EFFECT  OF  HEAT  UPON  THE  INSULATING  POWER  OF 

GLASS 

Support  at  both  ends  over  a  wing-top  burner,  a  piece  of 
glass  rod  1/4  to  3/8  inch  in  diameter.  Wind  two  pieces  of 
No.  20  nichrome  wire  tightly  about  the  rod  at  a  distance 
of  1  cm.  apart,  and  connect  them  to  a  110-volt  circuit 
through  a  low-reading  ammeter  and  a  lamp  bank.  Con- 
nect a  voltmeter  across  the  electrodes.  Heat  the  rod 
and  wires,  read  current  and  pressure  and  compute  the 
resistivity  of  the  glass.  Is  there  any  polarization? 

EXPERIMENT  19 
THE  ELECTROLYSIS  OF  GLASS  AT  300°  C. 

Prepare  about  300  .gr.  of  sodium  amalgam  by  alloy- 
ing three  percent  of  sodium  with  warm  mercury,  or  by  the 
electrolysis  of  a  saturated  salt  solution  with  a  mercury 
cathode.  Place  this  in  a  small  iron  crucible  and  support 
in  its  center  a  test-tube  containing  pure  mercury.  Weigh 
accurately  both  the  test-tube  and  the  mercury.  Connect 
the  crucible  as  anode  and  the  mercury  as  cathode  with 
a  rheostat,  ammeter  and  voltmeter;  heat  to  300°  C. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     27 

(pyrometer)  and  electrolyze  for  an  hotir  or  more.  Before 
closing  the  circuit,  record  the  E.M.F.  of  the  cell,  and  dur- 
ing electrolysis  open  the  circuit  from  time  to  time  and 
record  the  polarization. 

Reweigh  the  mercury  and  the  test-tube,  and  examine 
the  latter  for  any  sign  of  corrosion  or  chemical  action. 
Place  2  to  3  c.c.  distilled  water  on  the  mercury,  together 
with  a  piece  of  red  litmus  paper.  Inference?  The 
ampere-hours  passed,  multiplied  by  0.86  gives  the  weight 
of  sodium  in  grams  which  should  have  been  deposited  at 
100  percent  current  efficiency.  How  does  this  agree 
with  the  gain  in  weight  of  the  mercury? 

The  Internal  Resistance  of  Primary  and  Storage  Cells 

Let  E  =  the  E.M.F.  generated  by  the  cell; 

Let  R  =  its  internal  resistance,  and 

Let  R/  =  the  resistance  of  the  rheostat  connected  in 
series  with  the  cell,  and  assume  that  the  resistance  of  the 
connecting  wires  is  negligible  in  comparison,  as  is  usually 
the  case.  When  the  cell  delivers  the  current  I 

E  =  IR  +  IR'  or  IR'  =  E  -  IR. 

If  now  the  fall  of  potential  Er  across  the  rheostat  be  read, 
and  this  be  substituted  above, 

E  -  E' 
E'  =  E  -  IR  or  R  =  -  —=-• 

The  last  equation  is  the  one  usually  employed  in  measur- 
ing the  internal  resistance  of  primary  and  storage  cells. 
Note  that  for  exactness  this  equation  requires  that  E 
be  the  actual  E.M.F.  generated  by  the  cell  when  delivering 
the  current  I.  The  value  of  E  is  actually  found  by  read- 
ing the  E.M.F.  when  the  cell  is  delivering  zero  current,  i.e. 
on  open  circuit.  This  open  circuit  E.M.F.  might  be 
taken  immediately  before  or  immediately  after  the  cur- 
rent reading.  Which  will  represent  more  nearly  the 
E.M.F.  generated  by  the  cell  when  the  current  I  flows? 


28     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

EXPERIMENT  20 
THE  INTERNAL  RESISTANCE  OF  PRIMARY  CELLS 

Measure  the  internal  resistance  of  a  wet  cell,  also  of  a 
new  and  of  an  old  dry  cell.  For  measuring  the  resistance 
of  primary  cells,  it  is  best  to  draw  only  the  smallest  cur- 
rent which  will  give  a  sufficient  difference  between  E 
and  E'  to  be  read  accurately  on  the  instrument  used. 
Too  large  currents  cause  excessive  polarization. 

Connect  in  series  with  the  cell  to  be  tested  a  switch, 
a  low-reading  ammeter,  and  a  rheostat,  and  connect  a 
voltmeter  across  the  cell  terminals.  Close  the  circuit, 
read  the  current  and  E.M.F.,  open  the  switch  and  quickly 
read  E,  which  is  assumed  to  be  the  E.M.F.  generated  by 
the  cell.  Obtain  data  for  several  values  of  I,  and  com- 
pute the  resistance  by  the  equation  given  above. 

Is  it  clear  that  E'  is  the  fall  of  potential  over  the  rheo- 
stat, and  not  the  E.M.F.  of  the  cell  itself?  If  not, 
consult  the  laboratory  instructor. 

Since  the  voltmeter  takes  appreciable  current  when 
reading  the  open  circuit  E.M.F.,  the  assumption  that  this 
equals  the  voltage  generated  by  the  cell  is  incorrect. 
What  percent  of  the  true  E.M.F.  of  the  cell  is  E  in  the 
first  and  last  readings? 

Another  source  of  error  is  due  to  the  fact  that  your 
values  for  I  include  the  current  through  the  voltmeter. 
Compute  this  current  for  the  two  readings  just  referred 
to.  How  serious  is  this  error? 

Now  correct  the  first  and  last  values  of  resistance  for 
both  these  errors. 

EXPERIMENT  21 

THE  RESISTANCE  OF  STORAGE  CELLS 

Measure  the  resistance  of  small  storage  cells  of  both  the 
lead  and  the  nickel-iron  types.  Take  care  that  E  —  E'  is 
sufficiently  large  for  accuracy. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     29 

How  do  you  explain  the  difference  in  resistance  between 
the  lead  cell  and  the  wet  primary  cell? 

EXPERIMENT  22 

THE  RESISTIVITY  OF  SODIUM  PHOSPHATE 

By  the  fall  of  potential  method  using  110  volts  pressure, 
determine  the  specific  resistance  of  a  strong  solution  of 
sodium  phosphate,  first  with  lead,  then  with  aluminum 
electrodes. 

Is  the  difference  due  to  polarization?  Repeat  the 
measurement  using  one  lead  and  one  aluminum  electrode. 
Now  exchange  anode  and  cathode.  Keep  the  cell  for  the 
next  experiment. 

EXPERIMENT  23 

ELECTROLYSIS  WITH  ALTERNATING  CURRENT 

In  a  tumbler  electrolyze  a  solution  of  copper  sulphate 
with  carbon  electrodes  using  about  0.5  ampere  alter- 
nating current.  Result?  Now  put  the  cell  last  used  in 
experiment  22  in  series  in  the  circuit  and  pass  the  same 
current.  Explain. 

Faraday's  Law 

Faraday  said,  "The  chemical  decomposing  action  of  a 
current  is  constant  for  a  constant  quantity  of  electricity." 

The  principle  has  since  been  variously  stated:  "The 
weight  of  material  dissolved  or  deposited  at  either  elec- 
trode is  proportional  to  t  he  current,  to  the  time,  and  to 
the  chemical  equivalent  of  the  substance." 

"One  gram  equivalent  of  any  substance  is  dissolved, 
deposited,  or  decomposed  by  the  passage  of  96,540  cou- 
lombs (ampere-seconds)  of  electricity." 

For  practical  application,  the  most  useful  form  in 
which  to  memorize  this  law  is:  "For  every  26.8  ampere- 


30     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

hours,  1  gram  equivalent  of  any  substance  is  dissolved, 
deposited  or  decomposed." 

The  current  efficiency  of  all  electrolytic  processes  is 
determined  by  a  comparison  of  the  product  obtained  with 
the  yield  calculated  by  the  application  of  this  law.  In 
order  to  do  this,  it  is  necessary  to  measure  the  quantity  of 
electricity  in  coulombs,  or  ampere-hours,  and  the  amount 
of  the  desired  product  formed  in  the  same  time.  In 
technical  operations,  the  average  current  as  read  on  a 
calibrated  ammeter  may  be  multiplied  by  the  time  in 
hours,  or  the  same  result  may  be  read  directly  from  the 
dials  of  a  recording  ampere-hour  meter.  In  laboratory 
experiments  involving  small  currents,  it  is  customary  to 
place  in  series  with  the  electrolytic  cell  whose  efficiency  it 
is  desired  to  investigate,  a  coulombmeter  as  it  is  now 
called — formerly  known  as  a  voltameter.  For  measuring 
large  currents,  the  coulombmeter  may  be  placed  in  a 
shunt  circuit  through  which,  by  the  use  of  suitable  resist- 
ances, one-half,  one-tenth  or  any  desired  fraction  of  the 
total  current  is  made  to  pass. 

The  coulombmeter  is  an  electrolytic  cell  which  obeys 
Faraday's  law,  and  is  so  arranged  that  the  gas  or  metal 
deposited  can  be  accurately  measured  or  weighed. 

While  all  electrolytic  cells  obey  Faraday's  law  in  the 
sense  that  the  total  amount  of  material  deposited  at  the 
cathode  by  a  direct  current  is  strictly  in  accordance  with 
the  law,  some  of  the  metal  may  be  dissolved  by  an  acid 
or  other  corrosive  substance  present  in  the  electrolyte,  or 
some  other  substance  may  be  deposited  along  with  the 
one  desired  so  that  only  a  portion  of  the  current  is 
usefully  employed.  In  copper  refining,  the  large  amount 
of  sulphuric  acid  in  the  electrolyte  attacks  both  anode 
and  cathode,  with  the  result  that  the  current  efficiency 
at  the  cathode  is  below,  and  that  at  the  anode  is  above  100 
percent.  The  coulombmeter  has  this  advantage  over  the 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     31 

ammeter;  it  gives  the  ampere-hours  with  fair  accuracy 
in  spite  of  considerable  variation  of  the  current  and  it 
requires  no  attention;  while  with  an  ammeter,  the  cur- 
rent must  be  kept  constant  by  varying  the  resistance,  or 
the  variations  in  current  must  be  recorded,  which  neces- 
sitates watching. 

The  silver  coulombmeter  is  the  world's  standard  by 
means  of  which  the  value  of  the  ampere  is-  fixed.  For 
ordinary  laboratory  use,  however,  the  cheaper  and  less 
troublesome  copper  coulombmeter  is  sufficiently  accurate. 

EXPERIMENT  24 

THE  CONSTRUCTION  OF  A  COPPER  COULOMBMETER 

Oettel,  Practical  Exercises  in  Electrochemistry,  pp.  16, 
22.  The  electrolyte  should  consist  of  150  g.  of  copper 
sulphate,  50  g.  of  sulphuric  acid,  50  g.  of  alcohol  and 
1  liter  of  water.  A  rectangular  battery  jar  makes  a 
good  container.  Two  anodes  of  heavy  sheet  copper 
should  be  cut  with  projections  for  hanging  them  on  the 
top  of  the  jar.  The  single  cathode  of  thin  sheet  copper 
should  be  accurately  weighed.  For  tests  of  several  hours' 
duration,  the  electrolyte  should  be  stirred  by  a  jet  of 
hydrogen,  or  a  mechanical  stirrer.  With  good  circula- 
tion, 3  amperes  per  sq.  dm.  of  cathode  surface  may  be 
used.  At  the  end  of  the  test,  wash  the  cathode  with 
distilled  water,  rinse  with  alcohol  and  dry  quickly  over  a 
flame.  Oettel  gives  1.182  g.  of  copper  deposited  per 
ampere-hour,  corresponding  to  a  current  efficiency  of 
99.58  percent. 

EXPERIMENT  25 

THE  ELECTROLYSIS  OF  WATER  MADE  CONDUCTIVE  BY 
SODIUM  HYDROXIDE 

Fill  the  graduated  glass  apparatus  for  the  decompo- 
sition of  water  with  a  ten  percent  solution  of  purest  sodium 


32    A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

hydroxide,  and  connect  in  series  with  it  a  rheostat,  an 
accurate  ammeter  of  low  range  and  a  small  copper  cou- 
lombmeter  made  as  above.  Current  should  be  passed 
for  a  few  minutes  to  saturate  the  solution  with  the  gases 
before  any  gas  is  collected  for  measurement.  Hold  the 
current  constant  during  the  test  by  adjusting  the  rheo- 
stat, which  must  be  capable  of  very  fine  adjustment. 
Calculate  the  ampere-hours  from  the  weight  of  copper 
deposited,  and  determine  the  current  efficiency  of  hydro- 
gen and  oxygen  production,  and  the  error  of  the  ammeter. 
The  necessary  constants  for  hydrogen  and  oxygen  follow. 

HYDROGEN  OXYGEN 
1  liter  gas  at  0°  and  760  mm. 

weighs 0.089873  g.  1.42900  g. 

1  ampere-hour  liberates 0.03759  g.  0.2983  g. 

To  read  the  volume  of  gases  collected,  adjust  the  level 
of  liquid  in  the  open  tube  to  that  in  the  gas  tube,  read 
the  volume  and  temperature  of  each  gas,  and  the  height 
of  the  barometer  in  millimeters.  To  obtain  the  pressure 
under  which  the  gas  was  measured,  subtract  from  the 
height  of  the  barometer  the  tension  of  aqueous  vapor 
at  the  temperature  of  the  gas.  Change  the  gas  tempera- 
ture from  centigrade  to  the  absolute  scale  by  adding  273. 
The  gas  volume  may  be  reduced  to  the  standard,  0°  C. 
and  760  mm.  by  use  of  this  equation 

v°  =  v<  x  -     273        x  pressure 

absol.  temp-         760 

State  the  two  gas  laws  involved  in  the  above  equation; 
also  that  involved  in  the  subtraction  of  the  tension  of 
aqueous  vapor  from  the  height  of  the  barometer.  Is  an 
error  introduced  by  the  fact  that  the  gases  are  measured 
over  a  solution  of  caustic  soda  instead  of  over  pure  water? 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     33 

EXPERIMENT  26 

THE  ELECTROLYSIS  OF  WATER  MADE  CONDUCTIVE  BY 
SODIUM  NITRATE 

Using  the  decomposing  cell  of  the  previous  experiment 
in  series  with  an  accurate  ammeter,  electrolyze  a  ten  per- 
cent solution  of  sodium  nitrate  until  an  amount  of  gas 
suitable  for  measuring  has  been  collected.  Compute  the 
current  efficiency  as  before.  Explain  the  results. 

Potential  and  Electromotive  Force 

When  any  conductor  of  the  first  class  (metallic  con- 
ductor) is  dipped  into  a  conductor  of  the  second  class 
(electrolytic  conductor)  a  difference  of  potential  results. 
A  different  first-class  conductor  gives  a  different  potential. 
If,  then,  two  different  metals  be  simultaneously  immersed 
in  the  same  electrolyte,  there  should  be  a  difference  of 
potential  between  the  metals.  In  other  words,  the 
arrangement  constitutes  a  voltaic  cell.  • 

EXPERIMENT  27 

MEASUREMENT  OF  THE  ELECTROMOTIVE  FORCE  BETWEEN 
CONDUCTORS  OF  THE  FIRST  CLASS 

Prepare  clean  sheets  (3X5  inches)  of  the  following 
materials:  iron,  zinc,  copper,  lead,  brass,  aluminum, 
carbon  and  lead  peroxide.  The  last  may  be  a  strip  cut 
from  the  positive  plate  of  an  old  storage  cell,  and  freshly 
formed  by  using  it  as  anode  in  dilute  sulphuric  acid  for 
eight  to  ten  hours. 

Using  a  voltmeter  having  a  range  of  3  or  5  volts, 
measure  the  E.M.F.  between  copper  and  the  other 
materials  in  a  ten  percent  solution  of  sodium  sulphate, 
noting  in  each  case  which  electrode  is  attached  to  the 


34     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

positive  terminal  of  the  voltmeter.  To  avoid  injury  to 
the  voltmeter  by  an  E.M.F.  in  the  wrong  direction,  first 
test  the  direction  of  E.M.F.  in  each  combination  by 
barely  touching  for  an  instant  the  tip  of  the  electrode 
to  the  solution. 

Arrange  the  electrodes  in  the  order  of  their  potentials, 
calling  the  lowest  zero,  and  assigning  a  numerical  value 
to  each  of  the  others.  From  your  measurements  should 
copper  be  called  electro-positive  to  zinc,  or  vice  versa? 
The  current  entered  the  voltmeter  at  the  positive  ter- 
minal; in  which  direction,  then,  did  it  flow  through  the 
voltaic  cell? 

Also  measure  the  potential  of  the*  other  electrodes 
against  lead  peroxide,  and  arrange  a  similar  series  with 
lead  peroxide  as  zero.  Compare  the  values  of  a  com- 
bination of  the  same  two  electrodes  in  each  series. 

During  measurements,  note  changes  of  E.M.F.  due  to 
polarization  at  one  electrode. 

Single  Potentials 

Further  measurements  would  bring  out  even  more 
clearly  the  additive  nature  of  the  E.M.F.  of  voltaic 
cells,  that  the  E.M.F.  is  the  sum  of  the  potentials  of 
the  separate  electrodes. 

It  is  evident  that,  if  the  potential  of  some  particular 
electrode  be  chosen  as  the  standard  or  zero,  by  measur- 
ing the  E.M.F.  between  this  and  other  electrodes,  a 
scale  of  electrode  potentials  can  be  arranged  similar 
to  our  thermometer  scales  of  temperature  differences. 
This  has  been  done,  and  the  situation  in  regard  to  po- 
tential scales  in  use  today  is  strikingly  like  that  of  the 
thermometer  scales.  We  have  three  thermometer  scales 
in  use,  two  of  them  with  different  zero  points,  and  all  of 
them  requiring  negative  values  to  indicate  the  range  of 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY    35 

temperatures  encountered  in  the  world.  There  are 
two  full  potential  scales  in  use,  and  measurements  are 
occasionally  made  in  terms  of  other  incomplete  scales, 
used  only  for  special  purposes.  The  potential  scales 
are  also  like  our  temperature  scales  in  that  they  require 
negative  values  to  express  the  potentials  of  many 
substances. 

The  student  should  particularly  observe  that  a  negative 
value  for  the  single  potential  does  not,  of  itself,  denote 
anything  unusual  or  distinctive  about  that  metal,  but 
that  the  negative  sign  results  from  the  arbitrary  location 
of  the  zero  at  a  particular  place  in  the  scale.  Had  the 
zero  been  located  at  the  bottom,  all  potentials  would 
have  been  positive;  and  if  at  the  top,  all  potentials 
would  be  negative,  yet  the  order  of  the  series,  and  the 
physical,  chemical  and  electrochemical  properties  of 
the  electrodes  would  remain  absolutely  unchanged. 

Several  of  the  standard  electrodes  referred  to  have  been 
used  for  the  purpose  of  following  the  changes  in  potential 
occurring  at  the  plates  of  the  lead  storage  cell  during 
discharge  and  charge.  The  zinc  electrode  formerly 
used  for  this  purpose  has  now  been  superseded  by  cad- 
mium, on  account  of  its  more  constant  potential  and 
less  rapid  corrosion  by  the  battery  acid.  A  lead  peroxide 
electrode  has  also  proved  satisfactory  for  this  purpose, 
if  allowed  to  stand  twenty  hours  after  forming  before  it  is 
used.  The  method  of  using  these  is  to  connect  one 
terminal  of  the  voltmeter  to  the  standard  electrode 
and  the  other  to  that  plate  of  the  cell  whose  potential 
is  desired.  Polarization  caused  by  the  voltmeter  current 
is  negligible  in  this  case,  for  the  cadmium  is  always 
anode  so  that  the  only  effect  of  the  voltmeter  current 
is  to  dissolve  the  cadmium  at  the  rate  of  15  or  20  milli- 
grams per  hour.  Since  the  composition  of  the  electrode 
and  the  electrolyte  in  contact  with  it  remain  practically 


36     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

unchanged,  the  potential  is  constant.  The  plates  of 
even  a  small  storage  cell  expose  an  enormous  surface 
to  the  electrolyte  so  that  the  trifling  amount  of  hydrogen 
liberated  on  them  by  the  voltmeter  current  does  not 
affect  their  potential  appreciably.  (For  a  discussion 
of  Polarization,  see  page  15.) 

With  the  small  electrodes  which  it  is  frequently 
necessary  to  use  in  the  laboratory,  the  use  of  a  voltmeter 
causes  a  considerable  change  in  the  potential  at  one  or 
both  electrodes.  For  measuring  accurately  the  E.M.F. 
between  small  electrodes,  the  potentiometer  described  on 
page  37  is  far  more  satisfactory  than 'the  voltmeter. 

The  Normal  Calomel  Electrode 

The  standard  of  potential  most  generally  used  is  the 
normal  calomel  electrode.  A  convenient  form  of  this 
consists  of  a  wide  mouth  bottle  of  120  c.c.  capacity, 
with  a  layer  of  mercury  1  cm.  deep  in  the  bottom, 
on  which  is  the  same  depth  of  mercurous  chloride, 
previously  washed  with  normal  potassium  chloride 
solution  and  shaken  with  mercury  to  remove  any  mer- 
curic chloride.  The  bottle  is  nearly  filled  with  normal 
potassium  chloride  solution.  Connection  with  the 
mercury  is  made  by  a  platinum  wire  sealed  in  the  tip 
of  a  glass  tube  containing  a  little  mercury,  from  which 
a  copper  wire  makes  contact  with  the  external  circuit. 
Connection  with  the  electrolytic  cell,  the  electrode  of 
which  it  is  desired  to  test,  is  made  by  means  of  potas- 
sium chloride  solution  contained  in  a  glass  tube  which 
makes  contact  with  the  main  body  of  solution  in  the 
calomel  electrode.  To  this  glass  tube  is  attached  a  short 
rubber  tube  ending  in  another  glass  tube  bent  at  right 
angles  and  drawn  down  to  a  small  point.  A  plug  of 
filter  paper  or  asbestos  is  fitted  tightly  into  the  end  of 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     37 

the  glass  tube  to  hinder  diffusion  of  solutions  into  the 
normal  electrode.  The  bottle  should  be  closed  herme- 
tically (Why?)  by  a  three-hole  rubber  stopper,  the 
third  hole  of  which  carries  a  small  dropping  funnel  filled 
with  normal  potassium  chloride  solution,  by  means  of 
which  the  plug  of  filter  paper  is  washed  out  after  use. 

This  electrode  has  several  features  to  recommend  it. 
The  materials  are  cheap  and  readily  obtainable,  it  is 
easily  constructed  and  does  not  polarize  (change  poten- 
tial) with  the  passage  of  a  minute  current  in  either 
direction.  The  student  should  explain  its  freedom  from 
polarization  by  considering  the  chemical  changes  which 
occur  when  the  mercury  is  anode,  and  when  cathode. 
See  page  39  for  precautions  in  its  use. 

The  value  —0.56  volt  has  been  assigned  to  its  potential. 
The  scale  of  electrode  potentials  constructed  with  this 
as  a  basis  resembles  the  Fahrenheit  scale  of  temperatures 
in  that  no  constant  of  nature  happens  to  coincide  with 
zero  of  the  scale. 

On  account  of  the  high  resistance  of  the  calomel 
electrode,  a  voltmeter  cannot  be  used  for  measuring 
the  E.M.F.  between  this  and  other  electrodes.  For  this 
purpose  a  potentiometer  is  required. 

A  Simple  Potentiometer  and  its  Use 

A  two-dial  rheostat  consisting  of  nine  100-ohm  coils, 
and  ten  10-ohm  coils  makes  a  cheap  and  satisfactory 
potentiometer.  Connections  should  be  made  to  x  —  y, 
the  points  whose  E.M.F.  is  to  be  measured,  as  shown 
in  Fig.  9. 

The  middle  posts,  C,  D,  of  the  potentiometer,  to  which 
the  sliding  arms  are  attached,  are  to  be  connected  in 
series  with  the  short-circuit  key  K,  the  galvanometer  G, 
and  the  points  x  —  y.  A  source  of  E.M.F.  indicated  at 


38     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

B,  and  slightly  greater  than  that  to  be  measured,  is 
connected  across  the  outer  terminals  E,  F,  of  the  poten- 
tiometer. One  or  two  good  dry  cells  will  serve  for  this 
purpose  and  their  E.M.F.  should  be  measured  from 
time  to  time  during  use  by  an  accurate  voltmeter.  In- 
stead of  using  a  voltmeter,  the  unknown  E.M.F.  may 
be  compared  with  a  standard  cell  by  connecting  the 
latter  across  one  end  of  a  D.P.D.T.  switch,  the  circuit 


FIG.  9. 

x  —  y,  G,  K,  across  the  other  end,  and  the  binding  posts 
C,  D,  to  the  blades  of  the  switch. 

The  unknown  E.M.F.  must  be  connected  so  that  it 
opposes  B  as  shown  in  the  diagram.  The  dials  are  ad- 
justed until  the  galvanometer  indicates  no  current,  when 
the  unknown  E.M.F.  at  x  —  y  equals  the  fall  of  poten- 
tial between  D  and  C.  The  value  of  this  is  obtained  by 
multiplying  the  E.M.F.  of  B  by  the  reading  of  the  poten- 
tiometer expressed  as  thousandths.  The  third  decimal 
place  is  obtained  from  the  swings  to  the  right  and  left 
of  the  galvanometer  needle  as  one  of  the  10-ohm  coils 
is  introduced  and  cut  out  between  C  and  D. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     39 


EXPERIMENT  28 

THE  EFFECT  OF  THE  MATERIAL  OF  THE  ELECTRODE  UPON 
POTENTIAL 

In  a  normal  solution  of  sodium  or  potassium  sulphate, 
measure  the  E.M.F.  between  the  normal  calomel  elec- 
trode and  electrodes  of  cadmium,  carbon,  copper,  iron, 
lead,  lead  peroxide,  silver,  tin  and  zinc. 

Connect  the  calomel  electrode  at  y  and  the  metal  at  x. 
When  so  connected,  i.e.  the  wire  from  the  calomel  elec- 
trode leading  to  the  positive  terminal  of  the  potentiometer, 
the  reading  should  be  marked  positive,  but  if  it  is  found 
necessary  to  change  the  calomel  to  the  negative  side  in 
order  to  secure  a  balance,  the  potentiometer  reading 
should  be  marked  negative.  This  change  is  most  con- 
veniently made  by  a  pole-reversing  switch  in  the  battery 
circuit,  by  which  the  polarity  of  the  potentiometer  is 
reversed,  leaving  the  galvanometer  circuit  undisturbed. 
A  convenient  tabular  form  of  record  follows : 


Electrode 

Potentiometer 
reading 

Battery 
E.M.F. 

Cell  E.M.F. 

Electrode 
potential 

Nickel  .  .  . 

In  case  the  E.M.F.  varies  with  time,  this  variation  should 
be  recorded. 

Compute  the  electrode  potential  by  multiplying  the 
potentiometer  reading  by  the  E.M.F.  impressed  across  its 
terminals.  The  product  is  the  E.M.F.  of  the  cell  consist- 
ing of  the  metal  and  the  calomel  electrode.  To  this  add 
algebraically  the  potential  of  the  calomel  electrode 
(—  0.56)  and  the  result  is  the  potential  of  the  metallic 
electrode. 


40     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

Form  an  electrochemical  series  by  arranging  the 
electrodes  in  the  order  of  their  potentials. 

EXPERIMENT  29 

THE  EFFECT  OF  THE  COMPOSITION  OF  THE  ELECTROLYTE 
UPON  POTENTIAL 

Measure  the  potential  of  the  same  electrodes  in  a 
tenth  normal  solution  of  potassium  cyanide,  and  com- 
pare the  results  with  those  obtained  in  the  last 
experiment. 

EXPERIMENT  30 

THE  EFFECT  UPON  POTENTIAL  OF  CONCENTRATION  OF 
THE  ELECTROLYTE 

Measure  the  potentials  of  copper,  iron,  and  zinc  in  a 
saturated  solution  of  salt,  and  in  the  same  diluted  to  a 
tenth,  a  hundredth,  and  a  five-hundredth  of  the  original 
concentration. 

EXPERIMENT  31 

THE  EFFECT  OF  TEMPERATURE  UPON  ELECTRODE 
POTENTIAL 

Measure  the  potentials  of  copper,  iron,  and  zinc  in  a 
ten  percent  salt  solution  at  room  temperature,  at  about 
0°  C.,  and  at  100°  C. 

EXPERIMENT  32 

THE  EFFECT  OF  DISSOLVED  AIR  UPON  ELECTRODE 
POTENTIAL 

Measure  the  potential  of  platinum,  copper  and  zinc  in 
a  ten  percent  salt  solution  at  atmospheric  pressure,  then 
exhaust  the  air  by  means  of  the  laboratory  filter  pump, 
taking  readings  of  time  and  potential. 

For  doing  this,  put  the  electrolyte  into  a  bottle  similar 
to  the  calomel  electrode.  Drill  a  very  small  hole  through 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     41 

a  two-hole  rubber  stopper,  through  it  pass  a  number  14 
or  16  copper  wire  so  that  it  fits  air  tight,  and  attach  the 
metal  electrode  to  this.  Insert  the  tube  of  the  calomel 
electrode  in  one  hole,  and  in  the  other  put  a  glass  tube 
connected  to  one  leg  of  a  glass  Y  tube.  Remove  the 
dropping  funnel  from  the  calomel  electrode,  and  insert 
in  its  place  a  glass  tube  connected  to  the  other  leg  of  the 
Y  tube,  and  connect  the  Y  tube  to  the  filter  pump. 

May  not  the  potential  of  the  calomel  electrode  change 
as  a  result  of  removing  the  air  from  its  electrolyte,  and  so 
mask  any  change  at  the  other  electrode?  Test  this  by 
comparing  its  potential  with  another  calomel  electrode 
before  and  immediately  after  exhausting  the  air.  How 
prevent  re-absorption  of  air  by  the  calomel  electrode  be- 
fore its  potential  can  be  measured? 

EXPERIMENT  33 

THE   EFFECT  OF   GASES    OTHER  THAN  AIR   UPON 
ELECTRODE  POTENTIALO 

Measure  the  potentials  of  copper,  zinc  and  platinum  in 
a  ten  percent  salt  solution,  then  pass  in  a  stream  of 
illuminating  gas,  hydrogen  or  CO2,  and  continue  reading. 

The  Polarization  of  Voltaic  Cells 

The  polarization  of  a  voltaic  cell  may  be  readily  in- 
vestigated by  the  use  of  the  calomel  electrode. 

EXPERIMENT  34 
THE  POLARIZATION   OF  A  SIMPLE  VOLTAIC  CELL 

Hang  sheets  of  amalgamated  zinc  and  of  copper  in  five 
percent  sulphuric  acid  and  connect  them  through  a 
switch,  ammeter  and  a  resistance  of  5  to  10  ohms.  Con- 
nect a  voltmeter  of  about  1500  ohms  resistance  across  the 
plates,  and  set  up  a  calomel  electrode  to  read  the  potential 


42     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

of  the  copper  plate.  Read  E.M.F.  and  potential,  then 
close  the  switch,  and  take  readings  at  intervals  of  two  to 
three  minutes.  What  is  the  cause  of  the  change? 

PRECAUTION.  In  order  to  avoid  including  the  IR 
drop  through  the  electrolyte  in  the  potential  reading,  the 
tip  of  the  calomel  electrode  should  be  placed  as  close  to 
the  copper  as  possible,  and  behind  the  electrode,  i.e.  out 
of  the  line  of  current  flow.  The  other  alternative  is  to 
open  the  switch  a  fraction  of  a  second  before  depressing 
the  key  of  the  galvanometer.  Test  both  methods. 


FIG.  10. 

EXPERIMENT  36 
POLARIZATION  IN  A.LECLANCHE  TYPE  WET  CELL 

Connect  the  cell  in  series  with  a  rheostat,  ammeter  and 
switch,  put  the  tube  of  the  calomel  electrode  into  the 
electrolyte  of  the  cell  and  connect  with  the  potenti- 
ometer circuit  as  in  Fig.  10. 

Connect  a  voltmeter  (3  volt  range)  to  the  blades  of  a 
D.  P.  D.  T.  switch,  and  connect  one  end  of  the  switch  to 
the  wet  cell,  the  other  to  the  single  dry  cell  which  ener- 


A  LABOEATORY  COURSE  IN  ELECTROCHEMISTRY     43 

gizes  the  potentiometer,  so  that  all  E.M.Fs.  may  be  read 
on  the  same  instrument.  Read  the  open  E.M.F.  of  the 
cell  and  the  potential  of  its  cathode.  Close  the  switch  S, 
adjust  the  rheostat  for  1  ampere  current,  and  read 
E.M.F.  and  potential  at  five-minute  intervals. 

Electromotive  Force  of  Decomposition 

Experiments  1  to  7  consisted  of  a  qualitalive  study  of 
some  of  the  chemical  changes  which  occur  in  electrolysis. 
It  is  now  proposed  to  study  another  phase  of  electrolysis, 
viz.,  the  relation  between  current  and  impressed  E.M.F. 
as  the  latter  in  gradually  increased  in  magnitude. 

EXPERIMENT  36 
THE  E.M.F.  OF  DECOMPOSITION  OF  SODIUM  CHLORIDE 

The  apparatus  required  consists  of  a  voltmeter, 
(range  5  to  6  volts)  an  ammeter  (range  1  ampere)  a 


FIG.  11. 

slide- wire  rheostat  of  150  to  200  ohms  resistance,  a 
switch,  and  two  carbon  or  graphite  plates  of  8  to  10 
square  inches  area. 

Place  the  carbon  electrodes  in  a  ten  percent  salt  solu- 
tion, connect  the  fixed  terminals  of  the  rheostat  to  a  6- 
or  10-volt  circuit,  and  by  means  of  the  sliding  contact 
transfer  to  the  electrodes  increasing  fractions  of  the 


44     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

fall  of  potential  over  the  rheostat.  The  voltmeter 
should  be  connected  as  in  Fig.  11,  so  that  the  current 
traversing  it  is  not  registered  by  the  ammeter. 

Do  not  close  the  switch  until  ready  to  record  data. 
Adjust  the  rheostat  to  give  the  least  possible  E.M.P. 
across  the  cell,  and  read  the  E.M.F.  and  current.  In- 
crease the  E.M.F.  by  0.2-volt  intervals  until  the  limit  of 
the  voltmeter  or  ammeter  is  reached,  reading  the  current 
corresponding  to  each  E.M.F.  Does  the  current  corre- 
spond to  that  calculated  from  Ohm's  law?  Using  values 
of  E.M.F.  as  abscissae  and  of  current  as  ordinates,  plot 
the  current-E.M.F.  curve. 

The  results  are  characteristic  of  electrolysis  with  in- 
soluble electrodes.  The  E.M.F.  at  which  the  current 
rises  suddenly  is  called  the  E.M.F.  of  decomposition. 
What  is  its  value  in  this  case?  Some  have  attempted 
to  fix  upon  a  definite  value  by  extending  the  straight  part 
of  the  curve  backward  until  it  cuts  the  E.M.F.  axis, 
calling  the  point  of  intersection  the  E.M.F.  of  decom- 
position. Draw  such  a  line  to  your  curve. 

EXPERIMENT  37 
THE  COUNTER  E.M.F.  OF  POLARIZATION 

Why  does  Ohm's  law  not  apply  to  the  relation  of  cur- 
rent to  E.M.F.  in  the  last  experiment?  To  learn  this, 
repeat  the  experiment  opening  the  switch  after  reading 
the  current  at  each  value  of  E.M.F.,  and  as  quickly  as 
possible  read  the  counter  E.M.F.  of  polarization.  Plot 
the  current-E.M.F.  curve  as  before,  and  on  the  same 
sheet  plot  the  polarization-E.M.F.  curve,  marking 
values  of  polarization  as  ordinates  on  the  right  margin 
of  the  sheet. 

Why  does  an  impressed  E.M.F.  o  one  volt  produce 
no  current,  while  at  three  volts  current  passes?  Follow 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     45 

the  polarization  curve,  comparing  the  polarization  with 
the  E.M.F.  at  several  points.  Is  their  relation  a  con- 
stant one  throughout  the  curve?  Has  it  any  causal 
connection  with  the  current  curve? 


Measurement  of  the  Polarization  at  Anode  and 
Cathode 

We  have  seen  that  the  passage  of  current  through  an 
electrolytic  cell  may  develop  a  counter  E.M.F.  of  con- 


N    C    A 


FIG.   12. 


siderable  magnitude.  It  is  of  interest  to  learn  whether 
only  one,  or  both  electrodes  contribute  to  this  E.M.F. 
This  may  be  learned  by  introducing  a  metal  plate  which 
is  not  connected  to  the  electrolyzing  circuit,  and  measur- 
ing the  E.M.F.  between  this  and  each  of  the  active 
electrodes  at  the  outset,  and  for  each  value  of  the 
polarization.  The  connections  are  given  in  Fig.  12. 

S2  is  a  D.  P.  D.  T.  switch,  S3  a  S.  P.  D.  T.  switch,  R  a 
slide-wire  rheostat,  A  the  anode,  C  the  cathode,  and  N 


46     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

the  reference  electrode.  To  read  the  total  polarization 
throw  S2  to  the  right  and  then  open  the  line  switch  S 
for  an  instant.  For  the  E.M.F.  between  N  and  the 
cathode  throw  S2  to  the  left  and  S3  down;  for  the  anode 
throw  S3  up  while  S2  is  thrown  to  the  left.  By  com- 
paring the  E.M.F.  between  each  electrode  and  the 
standard  during  electrolysis,  with  its  initial  value,  the 
polarization  (change  of  potential  caused  by  the  passage 
of  current)  may  be  determined.  Note  that  in  voltaic 
cells  the  positive  terminal  of  the  voltmeter  is  always 
attached  to  the  cathode;  therefore  the  other  electrode 
is  the  more  electro-positive  of  the  two. 

EXPERIMENT  38 

THE   E.M.F  OF  DECOMPOSITION  AND  ANODE  AND 
CATHODE  POLARIZATION  FOR  ZINC  BROMIDE 

In  the  above  manner,  determine  the  E.M.F.  of  decom- 
position, the  total  polarization,  and  the  polarization  at 
each  electrode,  using  carbon  electrodes  in  a  solution  of 
zinc  bromide,  with  a  sheet  of  zinc  5X3  inches,  scoured 
with  pumice,  as  the  reference  electrode. 

It  will  be  noted  that  with  increase  of  impressed  E.M.F. 
the  anode  becomes  more  and  more  electro-negative, 
while  the  cathode  becomes  less  so. 

Plot  the  current-E.M.F.  curve  with  values  of  current 
at  the  right,  and  the  three  rjolarization  curves  (total, 
anode  and  cathode)  with  values  of  polarization  at  the 
left.  The  zero  of  polarization  should  be  placed  suffi- 
ciently high  so  that  all  negative  values  for  anode 
polarization  fall  upon  the  sheet. 

Instead  of  zinc,  carbon  might  have  been  used  as  a 
reference  electrode.  This  would  be  advantangeous  in 
that  the  polarization  of  each  electrode  could  be  read 
directly,  but  the  carbon  is  more  likely  to  change  its 
potential  during  the  experiment  than  the  zinc.  Why? 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     47 

EXPERIMENT  39 

THE   E.M.F.   OF  DECOMPOSITION   OF  SULPHURIC  ACID 
AND  THE  POLARIZATION  OF  LEAD  ELECTRODES 

By  the  method  of  the  previous  experiment,  determine 
the  E.M.F.  of  decomposition,  total  polarization,  and  the 
polarization  at  anode  and  cathode,  using  lead  electrodes 
in  normal  sulphuric  acid.  The  reference  electrode 
may  be  lead  peroxide  or  roughened  sheet  lead.  The 
former  has  the  advantage  of  greater  constancy.  If 
used,  it  should  be  prepared  as  follows:  Charge  a  strip 
from  the  positive  plate  of  a  lead  storage  cell  as  anode  for 
six  to  ten  hours,  until  oxygen  is  vigorously  evolved,  then 
use  as  cathode  at  a  low  current  for  twenty  minutes,  and 
allow  to  stand  over  night  in  dilute  sulphuric  acid  before 
use.  A  piece  of  the  negative  plate  from  a  storage  cell, 
charged  as  cathode,  makes  a  good  substitute  for  a  sheet 
lead  reference  electrode.  Plot  curves  as  in  experiment  37. 

E.M.F.  of  Decomposition  by  Observation  with  a  Lens 

Another  method  of  determining  the  E.M.F.  of  decom- 
position is  by  using  as  electrodes  platinum  wires,  sealed 
into  glass  tubes  with  5  to  6  mm.  exposed,  watching 
with  a  lens  for  the  first  product  of  decomposition  appear- 
ing at  either  electrode.  A  voltmeter  is  connected  across 
the  electrodes  to  measure  the  E.M.F.  applied,  which 
should  be  very  gradually  increased  until  decomposition 
is  attained.  The  E.M.F.  may  be  controlled  as  shown  in 
Fig.  12.  Several  trials  are  usually  necessary  to  find 
the  minimum  pressure  at  which  decomposition  occurs. 

EXPERIMENT  40 
THE   E.M.F.  OF   DECOMPOSITION  BY   OBSERVATION 

By  the  above  method  determine  the  E.M.F.  of  decom- 
position of  normal  sulphuric  acid,  of  a  normal  solution 


48     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

of  copper  sulphate,  and  of  fifteen  percent  nitric  acid.  Do 
products  appear  simultaneously  at  both  electrodes  in 
every  case?  If  not,  continue  to  raise  the  pressure  until 
deposition  occurs  at  the  other  electrode.  Decomposition 
must  have  occurred  when  the  first  product  was  liberated; 
why  then  did  not  the  product  appear  at  the  other  elec- 
trode also? 

The  use  of  the  voltmeter  for  measuring  potentials 
requires  electrodes  of  several  square  inches  area  in 
order  to  prevent  serious  alterations  of  potential  by 
the  current  required  to  operate  the  voltmeter.  Plati- 
num electrodes  of  this  size  are  often  unavailable,  in 
which  case  the  potentiometer  shoulct  be  substituted 
for  the  voltmeter. 

EXPERIMENT  41 

THE  MEASUREMENT  OF   POLARIZATION  BY  USE  OF  THE 
POTENTIOMETER 

With  platinum  electrodes,  a  potentiometer  and 
milliammeter,  determine  the  decomposition  and  the 
counter  E.M.F.  of  polarization  of  normal  hydrochloric 
ac.d,  hundredth  normal  hydrochloric  acid,  fifteen  percent 
nitric  acid,  and  normal  sulphuric  acid. 

Connect  as  in  Fig.  11  (page  43)  except  that  the 
potentiometer  (Fig.  9)  is  substituted  for  the  voltmeter, 
and  the  negative  terminal  of  the  latter  is  changed  to  the 
other  side  of  the  switch.  Why?  For  each  value  of 
E.M.F.  read  the  current,  then  open  the  switch  and  in- 
stantly tap  the  galvanometer  key.  If  a  special  combined 
switch-and-key  (described  on  page  52)  is  available,  the 
interval  between  opening  the  switch  and  closing  the 
galvanometer  circuit  may  be  shortened.  Note  the 
E.M.F.  at  which  products  of  decomposition  are  evolved 
at  the  electrodes. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     49 

Plot  the  current-E.M.F.  and  polarization-E.M.F. 
curves  as  usual.  Compare  the  E.M.F.  of  decomposition 
of  sulphuric  acid  with  the  value  obtained  in  experiments 
39  and  40.  Compare  your  results  with  Le  Blanc's  values 
in  Table  11  page  144. 

The  theory  of  electrolytic  dissociation  postulates 
that  the  molecules  of  electrolytes  are  decomposed  by 
the  mere  act  of  dissolving;  it  might  therefore  be  expected 
that  the  only  E.M.F.  required  to  send  current  through 
electrolytes  would  be  that  due  to  the  ohmic  resistance, 
and  that  conduction  in  all  electrolytes  would  at  all  times 
obey  Ohm's  law.  How  do  you  reconcile  the  theory  of 
electrolytic  dissociation  with  your  experiments?  In 
this  connection,  try  the  experiment  which  follows. 

EXPERIMENT  42 

DETERMINATION  OF  THE  E.M.F.  OF  DECOMPOSITION  OF 
NORMAL  COPPER  SULPHATE  WITH  COPPER  ELECTRODES 

This  may  be  carried  out  either  with  large  electrodes 
and  the  voltmeter  after  the  manner  of  experiment  36, 
or  with  small  electrodes  and  the  potentiometer  as  in 
experiment  41. 

EXPERIMENT  43 

THE  EFFECT  OF  INEQUALITY  IN  SIZE  OF  ELECTRODES 
UPON  E.M.F.  OF  DECOMPOSITION 

Using  platinum  electrodes,  one  not  over  a  centimeter 
square,  and  the  other  a  decimeter  square,  determine  the 
E.M.F.  of  decomposition  and  polarization  of  normal  sul- 
phuric acid  by  the  method  of  experiment  41.  Note  the 
lowest  E.M.F.  at  which  gas  appears  at  either  electrode. 
Explain. 


50     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

Further  Study  of  the  Polarization  of  Electrodes  and  Its 
Relation  to   the   Process   of   Electrolysis 

References : 

Richards  and  Landis,   Trans.  Amer.   Electrochem.  Soc., 

Vol.  4,  pp.  119-125. 
Watts,  Trans.   Amer.   Electrochem.  Soc.,   Vol.    19,   pp. 

91-106. 

Hitchcock,  Trans.  Amer.  Electrochem.  Soc,,  Vol.  25. 
Lehfeldt,  Electrochemistry,  p.  173. 

EXPERIMENT  44 

E.M.F.  OF  DECOMPOSITION  AND  POLARIZATION  AT  THE 
ELECTRODES  BY  THE  POTENTIOMETER 

With  electrodes  1  or  2  cm.  square,  a  voltmeter,  milli- 
ammeter,  potentiometer  and  calomel  electrode,  study 
the  electrolysis  of  the  following  solutions: 

a.  Normal  sulphuric  acid  with  platinum  electrodes. 

b.  Normal  hydrochloric  acid  with  platinum  electrodes. 

c.  Hundredth  normal  hydrochloric  acid  platinum  elec- 

trodes. 

d.  Fifteen  percent  nitric  acid. 

e.  Ten  percent   zinc    bromide  with  silver   electrodes. 
/.  Ten  percent  zinc  bromide  with  silver  cathode  and 

platinum  anode. 
If  possible,  carry  current  up  to  200  to  400  milliamperes. 

a.  Compare  values  with  those  obtained  in  experiment 
39,  and  explain  the  differences.     With  the  maximum 
E.M.F.  at  the  end  of  the  experiment,  open  the  line 
switch  and  follow  the  rate  of  diminution  of  polari- 
zation at  the  electrodes  by  readings  at  two-minute 
intervals  for  ten  to  fifteen  minutes. 

b.  and  c.  Why  the  difference  in  curves  of  anode  polariza- 
tion?    Compare  the  cathode  curve  with  that  in  a. 

d.  Why  does  not  hydrogen  from  nitric  acid  produce  the 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     51 

same  cathode  polarization  as  hydrogen  from  sul- 
phuric acid  in  a? 

e.  When  maximum  polarization  is  attained,  open  the 
line  switch  S6  and  follow  the  rate  of  depolarization  as 


FIG.  13. 


directed  in  a.     Explain  the  different  rates  of  de- 
polarization observed  in  a  and  c. 
The  electrical  connections  are  shown  in  Fig.  13. 


52     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

Although  this  looks  complicated,  it  is  easily  under- 
stood by  observing  that  it  consists  of  four  distinct 
circuits : 

1.  The   electrolyzing   circuit   comprising   a   source   of 
E.M.F.,  a  line  switch  S6,  a  rheostat  R,  milliammeter  MA, 
switch-key  K,  and  electrolytic  cell. 

2.  The  potentiometer  circuit,  consisting  of  battery  B, 
reversing  switch  S4,  potentiometer,  key  K,  galvanometer, 
cell,  and  switch  82. 

3.  Switches  S2  and  S3  for  substituting  in  the  potenti- 
ometer circuit  the  calomel  electrode  in  place  of  either 
anode  or  cathode. 

4.  Switch  S5  by  which  either  the  E.M.F.  of  the  electro- 
lytic cell  or  of  the  battery  B  may  be  read  on  the  volt- 
meter. 

The  key  K  is  a  combination  of  a  short-circuit  galvan- 
ometer key  (binding  posts  b,  c,  d,)  and  a  double-pole  line 
switch  (binding  posts  mn  and  m'n')  operated  by  a  single 
lever.  Normally  the  line  circuit  is  closed,  the  galva- 
nometer is  short-circuited,  and  the  potentiometer  circuit  is 
open.  On  depressing  the  key,  the  line  is  opened  and  then 
the  potentiometer  circuit  is  closed. 

It  may  require  three  hours  to  set  up  the  apparatus. 
Once  set  up,  it  should  be  left  in  position  for  use  by  other 
members  of  the  class. 

A  convenient  tabulation  for  data  and  results  follows: 


Potentiometer 
readings 

Volts,  polarization 

jj 

Mil- 
amp. 

!jf 

Q 

Cath- 
ode 

Anode 

1 

Cath- 
ode 

Anode 

Diff. 

0 

i 

0.0 

0 

4.169 

.010 

-.100 

-  .  090  .  042!  -  .  977  -    .  935'  .  042 

0.2        0 

4.169 

.050  -.100  -.150  .208  -.977-1.185  .208 

A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     53 

The  line  switch  S6  must  not  be  closed  until  after  tak- 
ing one  set  of  readings  to  find  the  initial  potential  of  each 


0.4 


2.4  Volts 


FIG.  14. 


electrode.  Increase  the  impressed  E.M.F.  by  steps  of  0.2 
volts.  Note  the  first  appearance  of  gas  at  each  electrode, 
and  watch  the  behavior^ of  the  milliammeter  for  any  infor- 


54     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

mation  which  this  may  give  in  regard  to  the  nature  of 
electrolysis.     A   comparison  of  the    difference    between 


-2.0 


anode  and  cathode  potentials  with  the  total  polarization 
gives  a  check  on  the  accuracy  of  the  measurements. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     55 

For  computing  the  polarization  from  the  potentiometer 
readings,  see  page  38. 

Plot  curves  of  current,  and  polarization  vs.  E.M.F.  as 
indicated  in  Fig.  14,  which  shows  the  electrolysis  of 
normal  sulphuric  acid  with  platinum  electrodes  1  cm. 
square,  placed  1  cm.  apart. 

In  your  experiments  does  the  polarization  of  a  plati- 
num cathode  by  hydrogen  follow  the  same  curve  in  dif- 
ferent electrolytes?  Explain.  What  have  the  various 
hydrogen-on-platinum  curves  in  common? 

Further  light  on  the  nature  of  electrolysis  may  be  ob- 
tained by  the  use  of  electrodes  of  unequal  size.  Fig. 
15  shows  three  separate  experiments  with  platinum 
electrodes,  one  of  1  cm.,  the  other  of  a  hundred  square 
centimeters  area,  a  centimeter  apart  in  normal  sulphuric 
acid. 

Not  only  is  the  polarization  at  anode  and  cathode 
dependent  on  the  relative  size  of  the  electrodes,  but  the 
total  polarization  curve  T,  and  the  current  curve  I  are 
also  considerably  modified. 

Experiment  1  was  made  with  a  cathode  1  cm.  square, 
and  an  anode  a  hundred  times  as  large.  The  first  point 
of  interest  is  that  hydrogen  is  evolved  at  an  E.M.F.  of 
1.1  volts,  which  is  0.6  volts  below  the  usual  value  for 
the  decomposition  of  water,  although  the  curve  for  the 
total  polarization  is  as  usual.  Note  that  anode  and 
cathode  polarize  about  equally  up  to  the  point  at  which 
hydrogen  is  first  evolved,  and  that  beyond  this  the  polar- 
ization of  the  anode  increases  rapidly. 

Before  starting  experiment  2,  the  electrodes  were 
removed  from  the  cell,  rinsed,  and  heated  to  redness  to 
expel  any  gases  that  might  have  been  absorbed  by  the 
platinum.  They  were  then  returned  to  the  same  elec- 
trolyte. Although  there  was  a  difference  of  potential  of 
1.57  volts  between  the  electrodes  when  they  were  re- 


56     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

moved  at  the  end  of  experiment  1,  they  are  seen  to  be  at 
the  same  potential  when  returned  at  the  beginning  of 
experiment  2,  which  was  merely  a  repetition  of  No.  1 .  The 
results  of  this,  so  far  as  the  polarization  of  the  separate 
electrodes  is  concerned,  are  quite  different  from  the 
first  trial.  The  polarization  of  the  cathode  increases 
much  more  rapidly  than  before,  while  the  potential  of 
the  anode  remains  almost  unchanged  up  to  an  impressed 
E.M.F.  of  0.8  volts.  Other  experiments  have  shown  that 
solutions  which  have  stood  long  enough  to  become 
saturated  with  air  have  a  marked  depolarizing  action  on 
the  cathode,  and  prevents  its  potential  from  rising  as  it 
would  do  if  the  dissolved  oxygen  were  not  present. 

In  experiment  3  the  smaller  electrode  is  used  as  anode, 
and  the  larger  as  cathode  in  a  fresh  lot  of  the  original 
solution  which,  of  course,  contained  dissolved  air.  The 
result  is  that  the  potential  of  the  large  cathode  remains 
practically  unchanged  up  to  an  impressed  E.M.F.  of  1.2 
volts,  when  oxygen  is  evolved  at  the  anode.  Beyond 
this  the  cathode  potential  rises  rapidly.  Notice  that 
hydrogen  was  obtained  in  only  two  of  the  three  experi- 
ments. Study  the  curves  and  predict  what  E.M.F. 
would  have  been  required  for  the  evolution  of  hydrogen 
in  experiment  3.  The  current  rose  steadily  beyond  0.8 
volts  E.M.F.  Explain  this  difference  from  No.  1  and 
No.  2.  The  curves  of  total  polarization  in  No.  1  and 
No.  2  coincide.  That  for  No.  3  coincides  with  the 
others  for  only  half  its  length.  Why  not  throughout? 

Discharge  Potentials 

The  potential  which  an  electrode  must  attain  in  order 
that  a  particular  substance  shall  be  evolved  or  deposited 
upon  it  has  been  termed  the  discharge  potential  of  that 
substance. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     57 

Examine  all  cathode  curves  or  tabulated  data  of 
previous  experiments  for  the  first  evolution  of  hydrogen. 
Prepare  a  table  by  recording  from  each  experiment  the 
cathode  material,  E.M.F.,  current  density  in  amperes 
per  sq.  dm.,  and  the  polarization.  What  seems  to  be 
the  determining  factor  in  the  evolution  of  hydrogen, 
the  electrolyte,  the  E.M.F.,  or  the  current  density? 

EXPERIMENT  45 
THE  DISCHARGE  POTENTIAL  OF  HYDROGEN  ON  PLATINUM 

Using  the  potentiometer  and  small  cathodes  of 
platinum,  copper,  mercury  and  carbon,  determine  with 
the  aid  of  a  lens  the  lowest  potential  for  the  appearance 
of  hydrogen  upon  the  cathode,  also  for  its  escape  as 
bubbles,  in  solutions  of  tenth  normal  sulphuric  acid 
and  tenth  normal  hydrochloric  acid. 

EXPERIMENT  46 

THE  DISCHARGE  POTENTIAL  OF  HYDROGEN  ON  SEVERAL 

METALS 

With  cathodes  of  platinum,  nickel  and  iron  determine 
the  same  in  tenth  normal  solutions  of  sodium  hydroxide, 
potassium  sulphate  and  sodium  chloride. 

EXPERIMENT  47 

THE  DISCHARGE  POTENTIAL  OF  CHLORINE 

By  a  similar  method  the  discharge  potentials  of  oxygen, 
chlorine  and  bromine  may  be  found  for  insoluble  anodes. 
Determine  the  discharge  potential  of  chlorine  with  plati- 
num electrodes  in  tenth  normal  solutions  of  hydrochloric 
acid  and  sodium  chloride.  After  the  latter  determina- 
tion, electrolyze  for  ten  minutes  with  a  brisk  evolution 
of  gas  and  the  formation  of  some  sodium  hypochlorite, 
and  repeat  the  determination.  Explain  the  changed 
result. 


58     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

Overvoltage 

In  experiments  45  and  46  it  was  seen  that  the  discharge 
potential  of  hydrogen  varied  with  cathodes  of  different 
materials.  The  discharge  of  hydrogen  has  been  found 
to  take  place  upon  platinized  platinum  at  a  lower  poten- 
tial than  on  any  other  metal.  The  difference  in  volts 
between  the  discharge  potential  of  hydrogen  on  plati- 
nized platinum  and  on  any  other  metal  is  called  the 
overvoltage  of  hydrogen  on  that  metal.  The  over- 
voltage  of  oxygen  at  insoluble  anodes  has  also  been 
investigated,  but  less  completely  than  that  of  hydrogen. 
Overvoltage  is  an  important  factor  in  determining  the 
corrosion  of  metals,  the  reducing  power  of  different 
cathodes,  and  the  oxidizing  power  of  anodes. 

EXPERIMENT  48 
THE  OVERVOLTAGE  OF  HYDROGEN 

In  tenth  normal  solutions  of  sulphuric  acid  and  sodium 
chloride,  measure  the  discharge  potential  of  hydrogen 
on  lead,  cadmium  and  tin  and  find  the  overvoltage  of 
hydrogen  on  these  metals  by  comparison  with  the  dis- 
charge potential  on  platinized  platinum.  With  the 
aid  of  a  lens,  find  the  lowest  potential  for  the  permanent 
clinging  of  bubbles  of  hydrogen  to  the  electrode. 

Rise  of  potential  with  increase  in  current  density  varies 
with  different  cathodes.  Also  find  the  potential  for  a 
current  density  of  1  ampere  per  sq.  dm.  Potential 
readings  should  be  made  on  open  circuit. 

The  Passive  State  of  Metals  an  Electrochemical 
Phenomenon 

When  iron  is  immersed  in  fuming  nitric  acid,  it  is  for 
a  time  rendered  immune  to  attack  by  dilute  nitric  or 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     59 

sulphuric  acid.  Several  other  metals  may  similarly  be 
rendered  passive  or  inert  to  chemicals  which  ordinarily 
attack  them.  The  discovery  of  the  passive  state  of 
metals  is  usually  attributed  to  Schoenbein  in  1836,  but 
instances  of  passivity  were  observed  at  much  earlier 
dates.  In  the  Philosophical  Transactions  for  1790, 
page  374,  Bergman  notes  that  silver  dissolved  in  strong 
red  nitric  acid  is  not  precipitated  by  iron.  In  the  same 
year  Kier  found  iron  dipped  in  fuming  nitric  acid 
insoluble  in  the  same  acid  of  ordinary  strength. 

EXPERIMENT  49 
THE  PASSIVE  STATE  OF  IRON 

a.  Dip  small  strips  of  cleaned  sheet  iron,  or  wire  nails 

freed  from  grease,  into  dilute  solutions  of  silver 
nitrate  and  copper  nitrate. 

b.  Immerse  clean  iron  in  fuming  nitric  acid,  remove 

and  at  once  dip  into  the  above  silver  and  copper 
solutions. 

c.  Measure  the  potential  of  cleaned  iron  in  fuming 

nitric  acid,  interposing  a  small  vessel  containing 
sodium  or  potassium  nitrate  solution  between 
the  nitric  acid  and  the  calomel  electrode  to  keep 
the  acid  from  the  latter.  It  is  sometimes  said 
that  iron  is  ennobled  by  dipping  in  fuming  nitric 
acid;  can  you  justify  this  statement? 

EXPERIMENT  50 

THE  PASSIVITY  OF  IRON  ANODES  IN  CERTAIN 
ELECTROLYTES 

By  grinding  or  pickling,  remove  the  scale  from  two 
strips  of  sheet  iron,  rinse,  and  use  as  electrodes  in  a  ten 
percent  solution  of  sodium  or  potassium  nitrate  contain- 


60     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

ing  a  few  drops  of  concentrated  nitric  acid.  Connect 
in  the  circuit  a  low-range  ammeter,  a  voltmeter,  and  a 
calomel  electrode.  Arranged  to  read  the  potential  of  the 
sheet  which  is  to  become  anode.  How  should  the  volt- 
meter be  connected?  Before  closing  the  line  switch, 
determine  the  potential  of  the  anode.  Close  the  switch, 
read  E.M.F.,  current,  and  polarization  as  the  E.M.F. 
is  increased  by  intervals  of  0.25  volts.  Compute  the 
current  density  for  each  value  of  current. 

EXPERIMENT  51 

PASSIVITY  vs.  POTENTIAL 

In  strong  solutions  of  potassium  chlorate  (faintly 
acidified  by  nitric  acid)  and  of  potassium  dichromate, 
measure  the  potentials  of  iron  and  cadmium  anodes  at 
current  densities  of  zero,  about  0.5  and  1  ampere  per 
sq.  dm. 

EXPERIMENT  52 

THE  PASSIVE  STATE  MAY  SOMETIMES  BE  UTILIZED  IN 
REMOVING  A  COATING  OF  ONE  METAL  FROM  ANOTHER2 

Number  and  weigh  accurately  several  cleaned  strips  of 
sheet  iron.  Electroplate  one  with  copper,  another  with 
brass,  and  a  third  with  nickel.  Weigh  again.  Connect 
these  as  anodes  with  iron  cathodes  in  three  cells  in  series, 
with  an  electrolyte  consisting  of  sodium  nitrate,  a  small 
amount  of  sodium  or  potassium  nitrite  and  a  little  nitric 
acid.  Pass  2  amperes  for  ten  to  fifteen  minutes.  Re- 
weigh.  What  is  the  cause  of  the  exception? 

The  Corrosion  of  Metals 

For  over  a  half  century  the  corrosion  of  metals  has  re- 
ceived much  attention  from  chemists,  engineers  and 

2  A  Practical  Utilization  of  the  Passive  State  of  Metals,  C.  F. 
Burgess,  Trans.  Am.  Electrochem.  Soc.,  Vol.  4,  pages  31-36. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     61 

technical  men  in  general.  In  later  years  the  corrosion 
and  preservation  of  steel  has  become  a  subject  of  vast 
economic  importance.  It  is  now  generally  admitted  that 
electrolysis  and  voltaic  action  play  an  important  part  in 
the  corrosion  of  iron,  other  metals  and  alloys. 

References : 

The  Corrosion  of  Iron  and  Steel — J.  N.  Friend. 

Corrosion  and  Preservation  of  Iron  and  Steel — Cushman 
and  Gardner. 

The  Corrosion  of  Iron  from  the  Electrochemical  Stand- 
point— C.  F.  Burgess,  Trans.  Am.  Electrochem.  Soc., 
Vol.  21,  pages  17-54. 

Effect  of  Substances  on  Corrosion  of  Iron  by  Sulphuric 
Acid — O.  P.  Watts,  Trans.  Am.  Electrochem.  Soc.,  Vol. 

v    21,  pages  337-353. 

EXPERIMENT  53 

THE  EFFECT  OF  CONTACT  WITH  OTHER  METALS  ON  THE 
CORROSION  OF  IRON 

Clean  a  sheet  of  iron,  cut  it  into  strips  1  X  10  cm., 
number,  and  weigh  them.  To  each  strip  fasten  in  the 
form  of  the  letter  V,  one  of  the  following :  carbon,  copper, 
lead,  tin  and  zinc  (scour  the  metals  with  pumice  first), 
and  place  the  combinations  inverted  in  vessels  containing 
a  ten  percent  solution  of  ammonium  chloride.  Leave 
several  days  until  there  is  marked  corrosion  of  some  of  the 
strips.  Note  time,  remove,  clean  off  rust,  (see  page  90) 
and  weigh.  Explain  results. 

EXPERIMENT  54 

THE  EFFECT  OF  CERTAIN  SALTS  ON  THE  CORROSION  OF 
IRON  BY  SULPHURIC  ACID 

Dilute  one  volume  of  concentrated  sulphuric  acid  by 
20  volumes  of  water,  put  200  c.c.  into  each  of  a  number  of 


62     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

tumblers,  place  them  in  a  thermostat  at  30°  C  or  some 
other  convenient  temperature,  and  cover  by  a  watch  glass. 
To  one  add  a  drop  or  two  of  a  solution  of  platinum 
chloride;  to  another,  1  g.  sodium  arsenate;  to  others, 
1  g.  copper  sulphate,  1  g.  tin  chloride,  1  g.  cadmium 
sulphate,  1  g.  silver  nitrate,  and  to  one  make  no  addition. 
Clean,  number  and  weigh  strips  of  rather  heavy  sheet 
iron  3X3  cm.  and  put  one  in  each  tumbler.  At  the 
end  of  twenty-four  hours  or  more,  remove  the  strips, 
brush  clean,  dry  and  weigh.  How  do  you  account  for  the 
various  results? 

EXPERIMENT  55 
ACCELERATED  CORROSION 

It  is  sometimes  desirable  to  hasten  the  corrosion  of  a 
metal,  in  which  case  the  principle  illustrated  in  the  last 
two  experiments  may  be  utilized. 

Prepare  dilute  sodium  amalgam  by  electrolyzing  a 
strong  salt  solution  in  a  broad  shallow  dish  with  a  mercury 
cathode  and  a  platinum  or  carbon  anode.  Divide  the 
amalgam  between  three  tumblers  half  filled  with  distilled 
water.  Into  one  drop  a  few  pieces  of  chromium,  into 
another  drop  a  carbon  rod.  Guess  at  the  relative  rates 
of  corrosion  of  the  sodium  in  the  three  cases. 

By  arranging  to  collect  the  hydrogen  in  graduated 
apparatus,  an  exact  determination  of  the  rates  may  be 
made.  Try  the  effect  of  a  few  drops  of  acid  in  the 
third  cell. 

How  would  iron,  copper  or  zinc,  act  in  place  of  the 
chrominum?  What  are  the  qualifications  of  a  satis- 
factory material  for  this  purpose?  What  may  be  the 
effect  of  metallic  impurities  in  the  mercury  of  commercial 
cells  for  the  electrolysis  of  salt? 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     63 

EXPERIMENT  56 

EFFICIENCY  OF  THE   DEPOSITION   OF  ALKALI   METALS 
FROM  VARIOUS  SOLUTIONS 

Operate  three  cells  in  series  with  equal  weights  of 
mercury  as  cathode,  with  electrolytes  of  three  normal 
solutions  of  sodium  and  potassium  chlorides  and  sodium 
sulphate.  Use  anodes  of  platinum  or  carbon  in  the 
first  two,  and  platinum  or  lead  in  the  last.  Electrolyze 
an  hour  with  occasional  stirring  of  the  mercury.  Wash 
the  amalgam  quickly,  decompose  in  equal  volumes  of 
distilled  water  by  the  use  of  chromium,  and  titrate  the 
the  alkali  formed.  What  was  the  strength  of  amalgam? 
The  current  efficiency?  What  reasons  can  you  give  for 
the  different  results? 

The  Electrolytic  Separation  of  Metals 

When  two  metals  are  in  the  same  solution,  it  is  often 
possible  to  pass  an  electric  current  through  the  solution 
under  such  conditions  that  only  one  metal  will  be 
deposited,  thus  effecting  a  separation  of  the  two. 

The  more  important  factors  contributing  to  the  success 
or  failure  of  this  process  are: 

1.  The  relative  position  of  the  two  metals  in  the  series 
of  single  potentials.     If  the  metals  are  far  apart,  they 
can  probably  be  separated,  but  if  of  about  the  same 
potential,    both    deposit    together,    though    there   is    a 
tendency  for  the  metal  of  lower  potential  to  be  deposited 
first. 

2.  The   nature   of   the   electrolyte   often   determines 
success  or  failure.     For  example,  in  sulphate  solutions 
the  potentials  of  zinc  and  copper  differ  by  a  volt,  and 
separation  is  easy,  but  in  cyanide  solution  their  poten- 
tials are  about  the  same,  and  an  alloy  is  deposited. 


64     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

3.  The  acidity,  alkalinity  or  neutrality  of  the  electro- 
lyte is  important.     Acidity  is  especially  so,  if  the  poten- 
tial of  hydrogen  lies  between  the  potentials  of  the  two 
metals  which  it  is  desired  to  separate.     The  addition 
of    acid    then    renders    separation    more    certain.     In 
alkaline   solutions,    the   magnitude   and   relative   order 
of  the  potentials  of  several  of  the  common  metals  are 
greatly  changed. 

4.  The  current  density  at  the  cathode  affects  the  re- 
sult, even  when  there  is  a  sufficient  difference  of  potential 
between  the  two  metals  to  render  a  separation  probable. 
So  long  as  there  is  enough  of  the  metal  of  lower  potential 
in   actual   contact   with  the   cathode  to   carry   all  the 
current,   that  metal   alone  is   deposited.     If,   however, 
metal  is  being  deposited  faster  than  diffusion  can  supply 
it,  enough  of  the  second  metal  will  be  deposited  to  make 
up  for  the  lack  of  the  first  metal,  unless  hydrogen  comes 
between  the  two,  in  which  case  that  may  be  deposited 
instead  of  the  second  metal. 

5.  Circulation.     The    function    of    circulation    is    to 
assist    diffusion    in    maintaining    in    contact    with    the 
cathode  a  sufficient  amount  of  the  first  metal  to  carry 
all  the  current. 

6.  Concentration   of   the   electrolyte.     It   is   evident 
from  4  and, 5  that  current  density,  circulation  and  con- 
centration are  interdependent  factors — that  any  change 
in  one  of  these  permits,  or  makes  necessary,  a  change 
in  one  of  the  others. 

It  should  be  noted  that  some  metals,  e.g.  lead,  man- 
ganese and  cobalt  may  be  deposited  partly  or  wholly 
at  the  anode  as  oxides. 

For  the  experiments  immediately  following,  prepare 
five  percent  solutions  of  the  sulphates  of  copper,  iron, 
nickel  and  zinc. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     65 

EXPERIMENT  57 

DEPOSITION   FROM  AN   ELECTROLYTE   CONTAINING 
SEVERAL  METALS 

Fill  a  tumbler  one-fourth  full  with  the  solution  of 
ferrous  sulphate.  With  an  iron  anode  and  a  brass 
cathode,  electrolyze  at  1  ampere  per  sq.  dm.  Result? 
Add  an  equal  volume  of  the  solution  of  zinc  sulphate 
and  continue  electrolysis.  Result?  Add  the  original 
volume  of  copper  sulphate  and  continue.  Result? 

EXPERIMENT  58 
DEPOSITION  FROM  ACID  ELECTROLYTES 

In  three  separate  vessels  put  solutions  of  the  sulphates 
of  copper,  iron,  and  zinc,  with  anodes  of  the  metal  in 
solution,  and  brass  cathodes;  electrolyze  in  series.  Add 
concentrated  sulphuric  acid  to  each,  drop  by  drop,  with 
stirring,  up  to  five  percent.  Note  the  results  from  time 
to  time.  Finally  put  in  new  cathodes,  or  reverse  the 
original  one,  to  bring  a  clean  portion  of  it  into  the 
electrolyte. 

With  the  aid  of  tables  of  electrode  potentials,  supple- 
mented by  your  own  exper  ments,  explain  the  results 
obtained  in  experiments  57-58. 

EXPERIMENT  59 
DEPOSITION  OF  AN  ALLOY  OF  NICKEL  AND  COPPER 

With  a  nickel  anode  and  a  copper  cathode,  electrolyze 
the  solution  of  nickel  ammonium  sulphate.  Then  add 
half  as  much  of  the  copper  sulphate  solution  and  con- 
tinue. Result?  To  half  the  above  mixture  add  am- 
monia until  most  of  the  precipitate  re-dissolves,  then  add 
a  solution  of  potassium  cyanide  until  the  blue  color 
disappears.  Electrolyze.  Result?  Explain. 


66     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

EXPERIMENT  60 
THE  DEPOSITION  OF  BRASS 

Mix  equal  volumes  of  two  percent  zinc  sulphate  and 
two  percent  copper  sulphate  solutions.  Electrolyze  with 
brass  anode  and  a  lead  cathode.  Result?  Now  slowly 
add  dry  sodium  carbonate  with  stirring,  until  evolution 
of  carbon  dioxide  ceases,  then  add  a  solution  of  potas- 
sium cyanide — (deadly  poison) — until  the  blue  color  dis- 
appears and  the  precipitate  is  almost,  but  not  quite, 
dissolved.  Electrolyze  again.  Explain. 

« 
Electrolytic  Analysis 

Chemical  analysis  by  the  electrolytic  method  is  now 
developed  to  such  an  extent  that  it  cannot  be  adequately 
treated  in  a  book  on  general  electrochemistry.  For 
information  on  this  subject,  the  student  should  consult 

E.  F.  Smith— Electro  Analysis. 

Neumann-Kershaw — Electrolytic  Methods  of  Analysis. 

Classen-Boltwood — Analysis  by  Electrolysis. 

One  of  the  greatest  improvements  in  this  method  of 
analysis  is  the  shortening  of  the  time  from  the  ten  or 
twelve  hours  formerly  required  to  fifteen  or  thirty 
minutes. 

It  is  evident  that  any  method  for  the  rapid  deposition 
of  metals  must  provide  for  vigorous  circulation  in  order 
to  bring  every  atom  of  metal  in  contact  with  the  cathode 
in  the  limited  time  allowed  for  the  process.  A  descrip- 
tion of  various  forms  of  apparatus  suitable  for  rapid 
electrolysis  may  be  found  in  the  5th  edition  of  Smith's 
Electro  Analysis  pages  39-67.  Since  1909  the  author  has 
used  a  cheap  and  convenient  arrangement,  consisting 
of  a  cylinder  of  platinum  gauze  5  cm.  in  diameter  and 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     67 

4  cm.  high,  weighing  with  the  two  conducting  wires 
9  g.,  and  a  motor-driven  centrifugal  pump  of  hard 
rubber  with  a  platinum  wire  anode  wound  spirally  about 
it.  The  strip  of  gauze  used  as  cathode  is  16  X  4  cm., 
7.75  meshes  per  sq.  cm.,  wire  0.116  mm.  diam.  (50  meshes 
per  square  inch,  wire  0.004  inches  diam.)  Its  actual  sur- 


FIG.  16. 

face,  calculated  by  the  formula  S  =  2?rdlb \/n,  in  which 
1  =  length  of  gauze,  b  =  breadth  and  n  =  number  of 
meshes  per  sq.  cm.,  is  13  sq.  cm.  The  maximum  cur- 
rent is  fixed  by  the  carrying  capacity  of  the  conducting 
wires.  A  recent  modification  consists  in  substituting  a 
gauze  anode  for  the  wire.  The  electrolyzing  vessel  may 


68     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

be  a  tall  type  beaker,  or  an  inverted  bottle  with  the 
bottom  removed  and  a  tube  passing  through  the  rubber 
stopper  for  drawing  off  the  electrolyte.  The  cathode 
should  be  held  centered  by  a  frame  of  light  glass  rod. 
The  stirring  is  very  effective,  since  there  is  circulation 
from  top  to  bottom  as  well  as  rotation  of  the  electro- 
lyte, and  the  jets  thrown  by  the  pump  pass  directly 
through  the  gauze  cathode.  The  apparatus  is  shown 
in  Fig.  16. 

EXPERIMENT  61 

THE  ELECTROLYTIC  DETERMINATION  OF  COPPER 

Weigh  accurately  about  1  g.  of  c.  p.»copper  sulphate, 
dissolve  it  in  150  c.c.  of  warm  distilled  water,  add  five 
drops  of  strong  nitric  acid  and  electrolyze  in  the  appara- 
tus just  described.  Test  for  complete  precipitation  by 
drawing  out  1  c.c.  of  the  solution,  making  alkaline  by 
ammonia  then  acidifying  by  acetic  acid,  and  adding  a 
few  drops  of  a  solution  of  potassium  ferrocyanide.  A 
brownish  color  or  precipitate  indicates  the  presence  of 
copper.  After  precipitation  is  complete,  without  in- 
terrupting the  current,  displace  the  electrolyte  by  water 
so  that  the  copper  shall  not  be  attacked  by  the  free 
acid  present;  or  the  cathode  may  be  removed  very  quickly 
after  interrupting  the  current,  plunged  into  water,  rinsed 
with  distilled  water,  then  with  absolute  alcohol,  and 
dried  by  hot  air. 

Note  the  E.M.F.,  current,  time  required,  speed  of  rota- 
tion and  appearance  of  the  deposit.  Too  large  a  current 
will  produce  a  dark  or  spongy  deposit.  Make  one  or 
two  check  analyses. 

EXPERIMENT  62 

THE  ELECTROLYTIC  DETERMINATION  OF  NICKEL 

Dissolve  about  1  g.  of  pulverized  nickel  ammonium 
sulphate,  or  0.7  g.  nickel  sulphate  in  120  c.c.  of  water, 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     69 

add  4  g.  ammonium  sulphate  and  30  c.c.  of  strong 
ammonia.  Electrolyze  as  above.  Test  for  complete 
precipitation  by  adding  a  solution  of  hydrogen  sulphide 
to  1  c.c.  of  the  electrolyte.  A  brown  color  indicates 
nickel.  The  nickel  must  be  entirely  removed  from  the 
platinum  by  the  use  of  warm  sulphuric  or  nitric  acid. 

Some  Electrochemical  Puzzles 

EXPERIMENT  63 
INTERMEDIATE  ELECTRODES 

Electrodes  which  carry  current  without  metallic 
connection  with  the  external  circuit  are  called  inter- 
mediate electrodes. 

In  each  end  of  an  oblong  rectangular  glass  vessel 
5  or  6  inches  long,  place  a  small  carbon  electrode  and 
between  these  suspend  horizontally  a  carbon  rod  or 
platinum  wire  2  1/2-3  inches  long,  with  its  ends  toward 
the  electrodes.  Fill  the  cell  with  an  eight  percent  solu- 
tion of  potassium  bromide  (or  sodium  chloride).  Is 
the  carbon  rod  a  better  or  a  poorer  conductor  than  the 
electrolyte  beside  it?  Predict  how  current  will  flow 
through  the  central  part  of  the  cell  where  it  has  a  choice 
of  passing  by  electrolytic  or  metallic  conduction. 

Connect  a  voltmeter  to  the  terminal  electrodes, 
apply  an  E.M.F.  increasing  by  increments  of  0.25  volt, 
and  note  results.  Explain. 

EXPERIMENT  64 
A  ROTATING  INTERMEDIATE  ELECTRODE 

(W.  D.  Bancroft,  Trans.  Am.  Electrochem.  Soc.,  Vol.  7. 
p.  171) 

Across  the  middle  of  a  rectangular  glass  jar  fit  a 
vertical  partition  consisting  of  a  pine  board  1  inch  thick. 


70     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

Boil  this  in  melted  paraffine  until  well  impregnated.  In 
the  middle  of  the  top  bore  all/4  inch  hole  to  a  depth  of  4 
inches,  and  fix  a  pivot  in  the  bottom  of  the  hole.  Drill  a 
hole  in  one  end  of  a  graphite  rod  (1  inch  diam.)  to  fit  the 
pivot,  fix  another  bearing  at  the  upper  end,  and  cut  a 
groove  near  the  top  to  carry  a  twine  belt,  so  that  the  rod 
can  be  rotated  about  a  vertical  axis  by  a  motor.  This 
leaves  the  graphite  rod  exposed  to  the  electrolyte  on  each 
side  of  the  board  diaphragm. 

Fill  the  cell  with  a  solution  of  100  g.  copper  sul- 
phate and  about  10  g.  sulphuric  acid  per  liter,  put 
electrodes  of  sheet  copper  in  the  ends  of  the  cell,  and  con- 
nect an  ammeter  (range  1  ampere),  a  Voltmeter,  and  a 
rheostat  as  in  Fig.  11,  page  43. 

Obtain  data  for  current-E.M.F.  curves,  first  with  the 
graphite  rod  stationary,  then  when  it  is  revolving.  Try 
different  speeds  of  revolution.  Explain. 

EXPERIMENT  65 
ALUMINUM  ELECTRODES 

Insert  weighed  aluminum  electrodes  in  a  gas  coulomb- 
meter,  fill  it  with  fifteen  percent  aluminum  chloride  solu- 
tion, and  with  an  accurate  ammeter,  electrolyze,  at  about 
0.2  ampere.  Collect,  measure,  and  identify  the  gas  at 
each  electrode.  Determine  the  current  efficiency  for 
metal  and  gas  at  each  electrode.  Explain  the  anomaly. 

Electroplating 

The  following  books  on  electroplating  should  be  avail- 
able for  consultation  by  students: 

Barclay  and  Hains worth — Electroplating. 
Field — Principles  of  Electro-deposition. 
Langbein-Brannt — Electro-deposition . 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     71 

Watt-Philip — Electroplating  and  refining. 
Trans.  Amer.  Electrochemical  Soc.,  Vol.  23. 

"The  Brass  World"  and  "The  Metal  Industry"  con- 
tain many  articles  concerning  commercial  electroplating, 
and  the  questions  and  answers  in  each  issue  reflect  troubles 
and  current  practice. 

Students  are  advised  to  read  the  following  articles: 

1.  Composition  of  Electroplating    Solutions — Keith,  Trans.  A. 
E.  S.,  3,  227-44 

2.  Chemistry     of      Electroplating — Bancroft,    Tr.     A.    E.    S., 
6,  27-43. 

3.  Physical  Character  of  Metal  Deposits — Burgess  and  Ham- 
buechen,  E.  and  M.  I.  1,  204-7. 

4.  Physical     Characteristics    of     Electro-deposited     Metals — 
Johnson,  E.  &  M.  I.,  1,  212-4. 

5.  Alloying  of  Metals  as  a  Factor  in  Electroplating — Kahlen- 
berg,  E.  &  M.  I.,  1,  201-2. 

6.  Tests    on    Elliptical    Anodes — Burgess    and    Hambuechen, 
E.  &  M.  I.,  1,  347-8. 

7.  Adhesion  of  Electrolytic  Metal  Deposits — Burgess  and  Ham- 
buechen, J.  Phys.  Chem.,  7,  409-15. 

8.  Electro-deposition  of  Nickel — Kern  and  Fabyan,  School  of 
Mines  Quarterly,  29,  342-70. 

9.  Deposition    of    Nickel — Johnson,  Tr.    A.    E.    S.,    3,    255-9. 

10.  Efficiency  of  the  Nickel  Plating  Tank— Brown,  Tr.  A.  E.  S., 
4,  83-99. 

11.  Injurious  Effect  of  Acid  Pickles  on  Steel — Burgess,  E.  & 
M.  I.,  4,  7-11. 

12.  Phenomena   of    Metal    Depositing— Betts,    Tr.    A.    E.    S., 
8,  63-79. 

13.  Function  of  Addition   Agents  in  Electrolytes — Kern,   Tr. 
A.  E.  S.,  15,  441-74. 

14.  Electro-deposition     of   Lead     from     Perchlorate     Baths — 
Mathers,  Tr.  A.  E.  S.,  17,  261-72. 

15.  Unsolved    Problems    in    Electroplating — Hogaboom,    Tr. 
A.  E.  S.,  19,  53-9. 

16.  A  Modern  Electroplating  Plant— M.  &  Ch.  E.,  8,  274-5. 


72     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

The  Composition  of  Plating  Solutions 

A  collection  of  the  various  solutions  that  have  been 
proposed  for  the  electro-deposition  of  the  more  common 
metals  may  be  found  in  several  papers  published  in  Vol. 
23  of  the  Transactions  of  the  American  Electrochemical 
Society.  A  perusal  of  these  papers  will  suggest  many 
interesting  experiments. 

The  composition  of  the  following  plating  solutions  is 
stated  in  grams  of  solids  for  1  liter  of  water.  Current 
densities  are  given  in  amperes  per  sq.  dm.,  and  resistivities 
in  ohms  per  centimeter  cube. 

1.  Brass  Bath — Roseleur's,  from  Pfanhauser's  Elek- 
troplattirung . 

Sodium  carbonate,  dry,  Na2C03  10  g. 

Cupric    acetate,    pulv.     Cu(C2H302)2'- 

H20  14  g. 

Sodium  bisulphite,  NaHS03  14  g. 

Zinc  chloride,  fused,  ZnCl2  14  g. 

Potassium  cyanide,  100  percent,  KCN40  g. 
Ammonium  chloride,  NH4C1  2  g. 

Current  density  0.3  ampere.  E.M.F.  2.7  volts.  Resis- 
tivity 13.6.  Specific  gravity  1.0545  (7  1/2°  Be). 
Current  yield  sixty-five  percent.  Deposit  in  one  hour, 
0.0041  mm. 

Prepare  the  solution  by  dissolving  the  sodium  salts  in 
400  c.c.  of  warm  water,  stir  the  copper  and  zinc  salts  with 
200  c.c.  of  water  and  slowly  stir  this  into  the  first  solution. 
Dissolve  the  cyanide  in  the  remainder  of  the  water,  and 
stir  into  the  other  portion  of  the  bath,  when  the  precipitate 
should  dissolve.  Add  the  ammonium  chloride  and  boil 
for  an  hour,  replacing  the  water  evaporated. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     73 

2.  Copper  Bath,  Acid. 

Copper  Sulphate,  CuS04-5H2O  200  g. 

Sulphuric  acid,  cone.  H2S04  30  g. 

Current  density  1  to  3  amperes.  Resistivity  9.3. 
Specific  gravity  1.1417  (18°  Be).  Current  yield  100 
percent. 

3.  Copper  Bath,  Alkaline. 

Sodium  sulphite,  Na2S03  20  g. 
Sodium  carbonate,  cryst., 

Na2C03-10H2O  20  g. 

Sodium  bisulphite,  NaHSO3  20  g. 

Cupric  acetate,  Cu(C2H3O2)2-H2O  20  g. 

Potassium  cyanide,  100  percent  KCN  20  g. 

Current  density  0.3  amperes.  E.M.F.  2.9  volts.  Re- 
sistivity 14.3.  Specific  gravity  1.0507  (7°  Be).  Current 
yield  seventy-one  percent.  Deposit  in  one  hour  0.0056 
mm.  Temperature  20°  C.  Make  up  as  directed  under 
bath  No.  1. 

4.  Gold  Bath  for  regular  gilding  on  all  metals. 

Sodium  carbonate,  dry  Na2C03  10  g. 
Gold  (as  double  chloride  of  gold 

and  ammonium)  (NH4)2  AuCle  2  g. 

Potassium  cyanide,  KCN  7  g. 

Current  density  0.1  amperes.  E.M.F.  2.8  volts.  Re- 
sistivity 44.  Specific  gravity  1.0175  (2  1/2°  Be).  Cur- 
rent yield  ninety-nine  percent.  Deposit  in  one  hour, 
0.00184  mm.  Temperature  20°  C.  Anodes  of  gold  one- 
thiid  the  area  of  the  cathode. 

5.  Gold  Bath  for  hot  gilding  of  articles  that  are  to 
be    only  partly    covered   with   go]d.     Cyanide    attacks 
enamel  or  lacquer. 


74     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

Potassium  ferrocyanide,  K4Fe(CN)6     15  g. 
Sodium  carbonate,  dry  Na2C03  15  g. 

Gold  chloride,  AuCl3  2.65  g. 

Current  density  0.1.  E.M.F.  2.1  vo^s.  Resistivity 
18.3.  Specific  gravity  1.0247  (3  1/2°  Be).  Current  yield 
ninety-five  percent.  Deposit  in  one  hour  0.00123  mm. 
Since  gold  anodes  are  insoluble,  carbon  anodes  may  be 
used.  Temperature  50°  C. 

6.  Iron  Bath. 

Ferrous  sulphate,  FeS04-7H2O  150  g. 

Feirous  chloride,  FeCV4H20  75  g. 

Ammonium  sulphate,  (NH4)2S(34       100  g. 

Cunent  density  1  ampere.  This  bath  is  suitable  for 
refining  iron  and  yields  good  deposits  an  inch  thick. 
At  20°  C.  the  deposit  is  hard  and  brittle,  but  electrolysis 
at  80°-90°  yields  a  softer  metal. 

7.  Lead  Bath  for  refining. 

Lead  (as  PbSiF6)  50  g.  to  80  g. 
Hydrofluor silicic  acid, 

H2SiF6  100  g.  to  150  g. 

Gelatine  0.5  g. 

Current  density  1.2  to  1.6  amperes.  For  plating,  the 
amount  of  free  acid  may  be  diminished  to  two  or  three 
percent. 

8.  Nickel  Bath  for  iron  and  steel. 

Nickel  ammonium  sulphate,  Ni(NH4)2 
(S04)2-6H20  75  g. 

Current  density  0.3  ampere.  E.M.F.  3.5  volts.  Resis- 
tivity 24.6.  Specific  gravity  1.0479  (6  1/2° Be).  Current 
yield  91.5  peicent.  Deposit  in  one  hour,  0.0034  mm. 
Cast  anodes  should  be  half  the  area  of  cathode. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     75 

9.  Nickel  Bath  for  brass  and  copper,  not  for  iron. 

Nickel  sulphate,  NiS04-7H2O  50  g. 

Ammonium  chloride,  NH4C1  25  g. 

Current  density  0.5  ampere.  E.M.F.  2.3  volts.  Resis- 
tivity 17.6.  Specific  gravity  1.0357  (5°  Be).  Current 
yield  95.5  percent.  Deposit  in  one  hour,  0.0059  mm. 
Cast  anodes  should  be  half  the  area  of  the  Cathode. 

10.  Nickel  Bath  for  pointed  objects  and  for  the  direct 
nickeling  of  zinc. 

Nickel  sulphate  40  g. 

Sodium  citrate  35  g. 

Current  density  0.27  ampere.  E.M.F.  3.6  volts.  Re- 
sistivity 5 1.7.  Specific  gravity  1.0394(51/2°  Be).  Current 
yield  ninety  percent.  Deposit  in  one  hour  0.00301  mm. 
Rolled  anodes  should  have  two  and  a  half  times  the  area 
of  the  cathode. 

11.  Nickel  Bath  for  thick  deposits. 

Nickel  sulphate,  NiS04-7H2O  50  g. 

Ammonium  tartrate  (neutral),  (NH4)2- 

C4H406  36  g. 

Tannin  0.25  g. 

Current  density  0.3  ampere. 

12.  Black  Nickel. 

Nickel  ammonium  sulphate  60  g. 

Ammonium  sulphocyanide  15  g. 

Zinc  sulphate,  cryst.  7  g. 

Use  nickel  anodes  three  to  four  times  the  surface  of  the 
cathode.  Current  density  0.05.  E.M.F.  0.5  volt.  The 
deposit  takes  on  any  metal  which  can  be  nickeled,  but 
is  best  over  white  nickel.  The  full  black  is  obtained 
only  on  polished  metal.  The  solution  must  be  kept 


76     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

strictly  neutral  by  addition  of  nickel  carbonate,  as  acid 
makes  the  deposit  gray  or  streaky,  and  alkali  causes 
brittleness  and  flaking  off. 

13.  Platinum  Bath— Roseleur's. 

For  thin  For  Thick 

Deposits  Deposits 

Pfanhauser 

Ammonium  phosphate  20  g.  100  g. 

Sodium  phosphate  100  g.  100  g. 
Platinum,  as  platinum 

chloride  2.3  g.  10  g. 

Current  density  1  to  2  amperes,  E.M.F.  3  to  4  volts. 

Dissolve  the  platinum  chloride  in  100  c.c.  of  water. 
Dissolve  the  ammonium  phosphate  in  200  c.c.  of  water 
and  add  to  the  solution  of  platinum  chloride.  Dissolve 
the  sodium  phosphate  in  700  c.c.  of  water  and  stir  it  into 
the  platinum  solution,  when  the  precipitate  previously 
formed  will  dissolve.  Boil  until  the  odor  of  ammonia 
has  disappeared  and  add  water  to  make  up  for  evapora- 
tion. The  bath  should  have  an  acid  reaction,  and  should 
be  used  hot. 

14.  Platinum  Bath— Bottger's. 

Citric  acid  105  g. 

Caustic  soda  to  neutralize 

Ammonium  platinic  chloride  from  1.58  g.  PtCl4 

Dissolve  the  citric  acid  in  400  c.c.  of  water,  neutralize  by 
caustic  soda,  and  to  the  boiling  solution  add  the  ammo- 
nium platinic  chloride  formed  by  dissolving  the  platinum 
chloride  in  a  small  amount  of  water  and  precipitating  it 
by  0.5  g.  of  ammonium  chloride  dissolved  in  a  few  c.c.  of 
water.  Make  up  to  1  liter  with  water. 

15.  Silver  Bath — for  heavy  plating. 

Silver  as  silver  cyanide  25  g. 

Potassium  cyanide  27  g. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     77 

Current  density  0.3  ampere.  E.M.F.  1.3  volts.  Resis- 
tivity 28.8.  Specific  gravity  1.0338  (4  3/4  Be).  Current 
yield  ninety-nine  percent.  Deposit  in  on^  hour  0.0114 
mm.  Area  of  anodes  equals  area  of  cathode. 

16.  Silver  Bath — for  ordinary  plating. 

Silver  as  silver  chloride  10  g. 

Potassium  cyanide,  100  percent  20  g. 

Current  density  0.3.  E.M.F.  1.5  volts.  Resistivity 
35.  Specific  gravity  1.0175  (2  1/2  Be).  Current  yield 
100  percent.  Deposit  in  one  hour  0.0115  mm. 

17.  Silver  Striking  Solution. 

Silver    as    silver    cyanide  4  g. 

Potassium    cyanide  100  g. 

18.  Brightener  for  Silver  Bath. 

Carbon  bisulphide,  CS2         45  g.  or  35.5.  c.c. 
Regular  silver  bath  1000  c.c. 

Shake  thoroughly  and  allow  to  stand  twenty-four  hours 
before  use.  Some  advise  adding  a  volume  of  ether  equal 
to  the  carbon  bisulphide.  For  use,  add  0.7  c.c.  of  the 
above  to  each  liter  of  the  silver  bath.  This  may  best 
be  done  by  putting  the  proper  amount  of  brightener  in 
a  large  bottle,  adding  a  liter  or  two  of  the  silver  bath  and 
shaking  until  a  uniform  solution  is  obtained.  This  is  to 
be  thoroughly  stirred  into  the  bath  in  the  plating  tank. 
Larger  amounts  of  brightener  give  a  dull  deposit,  and  an 
excess  spoils  the  deposit. 

19.  Amalgamating  Solution  or  Quick  Dip. 

Mercuric  oxide  (led)  6  g. 

Potassium  cyanide  100  g. 


78     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

20.  Tin  Bath — Roseleur's. 

Sodium  pyrophosphate,  Na4P207         40  g. 
Tin  chloride  (fused),  SnCl2  16  g. 

Tin  chloride  (cryst.),  SnCl2-2H2O  4  g. 

Current  density  0.3  ampere.  E.M.F.  2  volts.  Resis- 
tivity 40.2  Specific  gravity  1.0357  (5°  Be).  Current 
yield  ninety-nine  percent.  Deposit  in  one  hour  0.0059 
mm.  Anode  area  equal  to  cathode.  This  solution  may 
be  used  for  direct  deposition  on  copper,  brass,  bronze  or 
zinc;  but  iron  or  steel  must  be  coppered  first  or  given  a 
preliminary  coat  of  tin  by  an  immersion  bath.  The  tin 
anodes  do  not  corrode  satisfactory  and  tin  salts  must 
be  added  occasionally  to  maintain  a  sufficient  amount  of 
tin  in  solution. 

21.  Tin  Baths. 

a        b  c 

Caustic  soda,  NaOH       90  g.  120  g.  125  g. 
Tin  chloride,   cryst.   Sn- 

C12-2H2O  30  g.     30  g.     50  g. 

Sodium     hyposulphite, 

Na2S203-5H2O  15  g.     60  g.     75  g. 

Sodium  chloride,  NaCl    15  g. 

Use  hot  or  "cold  with  anodes  of  pure  tin. 

22.  Tin  Bath — for  contact  tinning: 

Cream   of   tartar,    KHC4H4O6, 

saturated  solution  1000  c.c. 

Tin  chloride,  SnCl2:2H20  20  g. 

Small  objects  of  copper  or  brass  may  be  given  a  thin 
coating  by  boiling  in  this  bath  in  contact  with  pieces  of 
granulated  tin. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     79 

23.  Tin  Bath  —  for  tinning  iron  by  immersion. 

Ammonium  alum,  (NH4)2A12-(SO4)4' 

24H2O  25  g. 

Tin  chloride,   fused,  SnCl2  2  g. 

A  bright  coating  is  produced  on  clean  iron  by  30  to  60 
seconds  immersion  in  the  boiling  solution. 

24.  Zinc  Bath. 

Zinc  sulphate,  ZnS047H2O  100  g. 

Ammonium  chloride,  NH4C1  25  g. 

Ammonium  citrate,  (NH^sCeHsO?      40  g. 


Current  density  0.5  to  1.0.  E.M.F.  1.1.  to  2.2.  Re- 
sistivity 15.9.  Specific  gravity  1.0781  (101/2°  Be). 
Current  yield  100  percent.  Deposit  per  ampere-hour 
0.0173  mm. 

25.  Zinc  Bath:—  (Metal  Industry,  1906,  85). 

Zinc  chloride  60  g. 

Ammonium  chloride  30  g. 

Hydrochloric  acid  4  g. 

Glycerine  4  g. 

Use  anodes  of  zinc  and  of  antimonial  lead  in  equal 
numbers.  The  bath  remains  clear  and  gives  a  fine  white 
deposit.  (Addition  of  ammonium  sulphate  would  assist 
in  protecting  the  lead  from  corrosion.) 

26.  Zinc  Bath—  Hansen-  Van  Winkle  Co. 

Zinc  sulphate,  ZnS04-7H2O  150  g. 

Aluminum  sulphate,  A12(S04)3-18H2O  50  g. 
Sulphuric  acid  1  g. 

Coloring  and  Oxidizing  Metals 

A  few  of  the  many  formulae  for  the  production  of 
various  colors  on  metals  are  given  below. 


80     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

27.  Black  on  Brass. 

Ammonia,  cone.  160  c.c. 

Water  160  c.c. 

Sodium  carbonate,  dry  10  g. 

Copper  carbonate  freshly  precipitated  until  excess 
remains  undissolved.  Heat  to  70°  C.  and  immerse  the 
brass  only  until  the  desired  color  is  obtained.  The  color 
is  made  more  permanent  by  immersion  in  a  hot  ten 
percent  solution  of  caustic  soda. 

28.  Bright  Black  on  Brass. 

Muriatic  acid,  HC1  ,     800  c.c. 

White  arsenic,  As203  200  g. 

Antimony  chloride,  SbCl3  120  g. 

This  works  best  hot,  and  no  water  should  be  added. 
Dip  the  brass  repeatedly  until  the  desired  color  is  ob- 
tained. The  color  will  stand  light  scratch-brushing. 
The  addition  of  100  g.  of  ferrous  sulphate  gives  a  bluish 
black.  The  solution  is  said  to  work  best  electrolytically 
with  steel  anodes. 

29.  Olive  Green  on  Solid  Brass. 

Copper  sulphate  80  g. 

Ammonium  chloride  20  g. 

Water  1000  c.c. 

Boil  the  objects  in  this  solution. 

30.  Antique  Green  on  Brass. 

Nickel  ammonium  sulphate  60  g. 

Sodium  hyposulphite,  Na2S203'5H2O  60  g. 

Water  1000  c.c. 

Heat  to  70°  C.  and  dip  the  objects.  It  gives  a  dark, 
slaty  green. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     81 

31.  Antique  Green  on  Brass  or  Copper. 

Acetic  acid  100  g. 

Ammonium  chloride  30  g. 

Sodium  chloride  10  g. 

Cream  of  tartar  10  g. 

Copper  acetate  10  g. 

Add  a  little  water  and  smear  over  the  brass.  Allow  to 
dry  for  twenty-four  to  forty-eight  hours,  and  relieve 
the  high  lights  with  a  brush  touched  to  beeswax. 

32.  Hardware  Green  on  Brass. 

Ferric  nitrate  Fe(N03)3  8  g. 

Sodium  hyposulphite  45  g. 

Water  1000  c.c. 

Heat  to  70°  C.  and  immerse  the  object  a  few  seconds. 

33.  Brown  on  Solid  Brass  or  Copper. 

Potassium  chlorate,  KC1O3  40  g. 

Nickel  sulphate,  NiSO4-7H2O  20  g. 

Copper  sulphate,  CuS04-5H20  180  g. 

Water  1000  c.c. 

Dip  the  article  in  pure  boiling  water,  then  boil  in  the 
above  solution,  rinse,  dry  and  buff  with  a  fine  brush. 

34.  Brilliant  Blue  on  Brass  or  Copper. 

Lead  acetate  15  g. 

Sodium  hyposulphite  25  g. 

Water  1000  c.c. 
Heat  to  80°  C.,  immerse  the  cleaned  object  for  two 

to  twenty  seconds,  rinse,  dry,  and  lacquer.     The  color 
fades  after  several  months. 

35.  Brown  to  Black  on  Copper. 

Potassium  sulphide,  K2S  6  g. 

Ammonia,  cone.  5  c.c. 

Water  1000  c.c. 
Use  cold,  or  only  slightly  warm. 


82     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

36.  Deep  Black  on  Copper. 

Copper  nitrate  100  g. 

Water  300  c.c. 

Immerse  the  articles  or  paint  on  the  solution,  heat  on 
a  hot  plate  or  over  a  flame  to  convert  the  nitrate  to 
oxide.  If  desired,  this  treatment  may  be  followed  by 
immersion  in 

Potassium  sulphide  100  g. 

Water  1000  c.c. 

Hydrochloric  acid  10  c.c. 

37.  Black  on  Iron  or  Nickel. 

Lead  nitrate  90  g. 

Ammonium  nitrate  60  g. 

Water  1000  c.c. 

Heat  the  bath  to  60°  C.,  suspend  the  cleaned  articles 
as  anode  by  iron  wire,  with  lead  cathodes,  and  electrolyze 
until  black.  Remove,  rinse  in  hot  water,  dry  in  sawdust, 
and  buff  or  scratch-brush  according  to  the  finish  desired. 

38.  Black  on  Silver  or  "Oxidized  Silver." 

Ammonium  carbonate  12  g. 

Potassium  sulphide  6  g. 

Water  1000  c.c. 

Heat  to  80°  C.  and  immerse  the  articles  five  to  thirty 
seconds.  The  deposit  will  stand  scratch-brushing.  A 
one  percent  solution  of  barium  sulphide  acts  more  slowly. 

39.  Metallochromes. — Brilliant  multi-colored  deposits 
of  lead  pei  oxide  may  be  produced  upon  polished  nickel 
01  steel  used  as  anode  in  solutions  of  several  lead  salts. 

Caustic  soda,  NaOH  50  g. 

Lead  nitrate  or  acetate  5  g. 

Water  1000  c.c. 

Dissolve  the  two  solids  separately  in  small  amounts 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     83 


of  water,  pour  the  lead  solution  into  the  other,  then  dilute 
to  the  proper  volume. 

For  ready  leference,  data  in  regard  to  the  foregoing 
baths  have  been  collected  in  the  table  which  follows. 

TABLE  1 


No. 
of 
bath 

Metal 

Bath 

Metal  per  liter 

Per- 
cent 
effi-. 
ciency 

E. 
M. 
F. 

Curren 
sity 
amper 

sq. 
dm. 

t  den- 
in 
33  per 

Grams 

Gram 
equiva- 
lents 

sq. 
foot 

1 

Brass        Cyanide 

12.8      0.16 

65-70 

2.7 

0.3j  2.8 

2 

Copper 

Sulphate 

48.4 

1.58 

98 

3.2 

30.0 

3 

Copper 

Cyanide 

6.3 

0.10 

81 

2.9 

0.3 

2.8 

4 

Gold 

Cyanide 

2.0 

0.03 

95 

2.9 

0.1 

0.9 

6 

Iron 

Sulphate- 
chloride 

51.3 

1.83 

98 

1.0 

1.1 

10.0 

7 

Lead 

Silicofluor- 
ide 

\    80.0 
J  100.0 

fO.77 
10.97 

|1.6 
13.0 

28.0 

8 

Nickel 

Sulphate 

11.2 

0.38 

92 

3.5 
3-4 

0.3 

2.8 

13 

Platinum 

Pyrophos- 
phate 

f      2.3 
I    10.0 

fO.05 

10.20 

1.0 

9.3 

15 

Silver 

Cyanide 

2.5 

0.23 

99 

1,3 

0.3 

2.8 

20 

Tin 

Pyrophos- 
phate 

12.2 

0.20 

99 

2 

0.2 

1.8 

21 

Tin 

Hyposul- 
phite 

f    15.8 
I    26.3 

fO.26 
10.44 

24 

Zinc 

Sulphate- 
citrate 

22.7 

0.69 

100 

2.2 

1.0 

9.3 

26 

Zinc 

Aluminum 
sulphate 

34.1 

1.04 

100 

1.6 

15.0 

84     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

In  comparing  the  metal  content  of  the  different  baths 
with  the  current  densities  which  experience  has  shown 
may  be  safely  used,  the  controlling  factor  is  gram 
equivalents,  not  the  total  weight  of  metal  in  the  solution. 

It  is  expected  that  students  will  consult  this  table  to 
learn  the  current  densities  that  should  be  used  for  the 
experiments  in  plating. 

It  is  the  desire  of  the  electroplater  to  obtain  smooth, 
solid,  tough  and  adherent  deposits  of  metal.  This  is 
by  no  means  easy  to  secure  in  all  cases.  Important 
factors  influencing  the  nature  of  electrolytic  deposits  are : 

Principles  of  Electrodeposition 

1.  The  metal  deposited. 

2.  The  metal  receiving  the  deposit. 

3.  The  chemical  composition  of  the  electrolyte. 

4.  Gases  dissolved  in  or  evolved  from  the  electrolyte. 

5.  Insoluble  impurities. 

6.  Temperature. 

7.  Current  density. 

8.  Concentration  and  circulation  of  the  electrolyte. 

9.  Thickness  of  the  deposit. 

10.  Extent  of  anode  surface  and  ai  rangement  of  anodes. 

That  certain  metals  have  a  characteristic  foim  of  de- 
posit is  illustrated  by  the  spangles  in  which  lead  plates 
out  of  a  solution  of  its  nitrate  or  acetate,  and  the  needles 
which  silver  forms  when  deposited  from  a  solution  of  its 
nitrate. 

The  difficulty  of  securing  an  adherent  plating  upon 
highly  electro-positive  metals,  like  mangnesium  or 
aluminum,  is  well  known.  Even  zinc  is  sufficiently 
electro-positive  to  cause  trouble  in  plating  upon  it  from 
many  solutions. 

The  chemical  composition  of  the  electrolyte  has  long 
been  recognized  as  one  of  the  most  important  factors. 
This  is  attested  by  the  host  of  formulas  extant  for  the 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     85 

deposition  of  a  single  metal,  e.g.,  nickel.  Very  minute 
quantities  of  material  may  profoundly  affect  the  nature 
of  the  deposit,  as  recited  in  the  many  articles  upon  the 
use  of  addition  agents,  as  such  substances  are  called. 
Either  with  or  without  the  use  of  addition  agents  it  is 
often  important  to  use  a  certain  salt  of  the  metal  in  the 
electrolyte.  A  change  in  the  electrolyte  from  an  acid 
to  an  alkaline  reaction,  or  vice  versa,  may  spoil  some 
metal  deposits. 

Gas  bubbles  sometimes  cling  to  the  cathode  and  cause 
pits  as  the  metal  around  them  grows  in  thickness.  This 
is  especially  likely  to  occur  in  nickel  plating.  Another 
effect  of  gases  is  seen  in  the  hardness  and  brittleness  of 
electro-deposited  iron  and  nickel,  which  is  generally 
attributed  to  the  absorption  of  hydrogen  by  the  metal. 

A  solution  free  from  suspended  matter  is  one  of  the 
requisites  for  good  plating  from  stationary  solutions, 
but  in  a  few  cases  of  deposition  on  rotating  cathodes, 
such  material  is  added  in  large  amount  to  produce  a 
polishing  effect  on  the  deposited  metal.  It  is  the  common 
practice  to  stir  the  plating  solutions  only  at  the  close 
of  the  day's  work.  This  removes  the  layer  of  dense 
solution  from  the  bottom  of  the  tank  and  allows  time 
for  the  sediment  to  settle  again  before  the  bath  is  used. 

Temperature,  current  denity,  concentration  and  circu- 
lation are  interdependent  factors.  The  maximum  cur- 
rent density  allowable  without  spoiling  the  deposit 
seems  to  depend  on  the  maintainance  of  enough  dis- 
solved metal  in  actual  contact  with  the  cathode  to  carry 
whatever  current  is  passing.  The  role  of  concentration 
and  of  stirring  is  apparent.  Rise  of  temperature 
increases  the  rate  of  diffusion  of  dissolved  substances 
and  so  raises  the  maximum  current  density  permissible 
in  stationary  solutions.  It  has  long  been  customary  to 
use  gold  and  platinum  solutions  hot.  A  glance  at 


86     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

Table  1  of  page  83  shows  that  as  usually  used,  these 
baths  stand  out  from  the  others  as  notably  deficient 
in  metal.  The  reason  for  the  better  results  attained  in 
the  deposition  of  these  two  metals  from  hot  rather  than 
from  cold  solutions  is  obvious.  Many  platers  are  now 
reporting  more  satisfactory  results  from  a  hot  than  from 
a  cold  brass  solution.  Besides  improving  circulation, 
elevation  of  temperature  greatly  stimulates  corrosion  of 
the  anode,  and  so  lessens  a  recognized  difficulty  in  the 
operation  of  the  brass  bath,  viz.,  poor  anode  corrosion. 
Still  another  effect  of  high  temperatures  is  to  lessen  the 
absorption  of  hydrogen  by  nickel  and  iron,  and  so  permit 
the  production  of  softer  and  less  brittte  deposits  of  these 
metals.  As  the  advantages  of  hot  solutions  become 
recognized  by  platers,  they  will  be  used  far  more  generally 
than  at  present,  in  spite  of  the  trouble  and  cost  of  heating 
them. 

The  stirring  of  the  solutions  used  in  electrolytic 
analysis  has  cut  down  the  time  required  for  the  deposition 
of  a  given  weight  of  metal  from  about  eight  hours  to  fifteen 
minutes,  or  even  less,  but  the  stirring  of  plating  baths 
during  use  has  met  with  little  favor  on  account  of  the 
sediment  always  present  in  platers'  solutions.  In  the 
deposition  of  heavy  coatings  of  silver,  the  importance 
of  breaking  up  the  film  of  impoverished  solution  which 
forms  around  the  cathode  was  long  ago  recognized,  and 
various  mechanisms  have  been  used  for  moving  the 
objects  during  plating.  Similar  devices  are  now  coming 
into  use  for  the  rapid  production  of  heavy  nickel  deposits. 

The  advantages  of  a  concentrated  solution  have 
recently  been  recognized  in  nickel  plating,  and  the 
market  is  flooded  with  " high-power"  nickel  salts,  sold 
under  fancy  names,  which  in  their  essence  are  only  more 
soluble  salts  than  the  sparingly  soluble  double  sulphate 
of  nickel  and  ammonium  which  has  remained  the 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     87 

standard  for  so  many  years.  The  next  logical  step  would 
be  the  use  of  a  hot  nickel  solution. 

It  is  comparatively  easy  to  obtain  fairly  smooth  de- 
posits of  most  metals  for  several  minutes,  but  with  con- 
tinued electrolysis,  as  the  hours,  days  and  weeks  (in 
refining)  pass,  and  the  deposit  becomes  thicker,  it  loses 
its  original  smoothness,  and  continually  grows  rougher 
and  more  nodular,  so  that  it  is  an  exceptional  electrolyte 
that  yields  a  smooth  deposit  having  a  thickness  exceeding 
a  half  inch.  In  some  plating  solutions,  the  deposit  on 
a  polished  object  becomes  dull  from  incipient  roughness 
in  a  minute,  or  less,  while  in  other  baths  it  remains 
bright  for  four  or  five  minutes,  even  though  the  metal  is 
deposited  at  the  same  rate  in  both  cases. 

Extent  of  anode  surface  is  important  in  maintaining 
a  plating  bath  in  its  initial  condition  as  regards  acidity, 
neutrality,  etc.  As  the  current  density  at  the  anode 
increases,  there  is  a  tendency  for  the  efficiency  of  corro- 
sion to  diminish.  When  the  current  efficiency  at  the 
anode  is  less  than  that  at  the  cathode,  there  is  a  produc- 
tion of  free  acid,  free  cyanide,  etc.,  in  the  solution,  and 
when  the  efficiency  at  the  anode  exceeds  that  at  the 
cathode,  the  solution  tends  toward  alkalinity.  It  is 
evident  that  by  suitable  proportioning  of  the  extent  of 
anode  and  cathode  surfaces,  it  should  be  possible  to 
maintain  the  initial  condition  of  the  solution. 

Since  a  deposit  of  uniform  thickness  is  usually  desired, 
the  anodes  should  be  arranged  so  as  to  distribute  the 
current  uniformly  over  the  cathode.  With  pointed  ob- 
jects and  those  of  irregular  shape,  this  is  frequently  a 
matter  of  considerable  difficulty  and  sometimes  re- 
quires special  anodes,  made  to  conform  to  the  shape  to 
the  objects. 

An  ideal  plating  bath  would  possess  the  following 
properties : 


88     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

1.  A  current  efficiency  of  100  percent  at  both  anode  and 
cathode. 

2.  Simple  composition,  containing  only  compounds  of 
the  metal  which  is  to  be  deposited. 

3.  Very  soluble  metallic  salts,  so  as  to  permit  the  use 
of  a  high  current  density,  and  a  corresponding  saving  in 
size  of  the  plating  equipment. 

4.  No  oxidation  or  reduction  of  the  dissolved  salt  as  a 
result  of  the  passage  of  current.     Because  of  reduction, 
nitrates  are  unsuitable  as  electrolytes  either  for  plating 
or  refining. 

5.  Low  resistance  of  the  solution  in  order  that  large 
currents  may  be  used  with  an  E.M.F.  bf  6  volts  or  less. 

6.  Free  solution  of  the  anode  under  the  influence  of  the 
current  without  the  formation  of  a  film  of  oxide  or  other 
compound  upon  it.     The  anode  should,  however,  not  be 
attacked  when  the  bath  is  idle. 

Since  a  current  efficiency  of  100  percent  at  both  elec- 
trodes is  usually  unattainable,  the  next  best  condition  is  to 
have  equal  efficiencies  at  both  electrodes,  in  order  that 
the  metal  content  of  the  bath  may  remain  unchanged 
with  use.  There  are  a  few  cases,  e.g.  platinum  baths,  in 
which  it  is  not  possible  to  use  soluble  anodes,  and  metal 
must  be  added  from  time  to  time  in  the  form  of  soluble 
salts. 

While  simplicity  of  composition  is  desirable,  it  is 
often  necessary  to  add  foreign  substances  to  overcome 
defects  in  the  operation  of  the  bath.  Acids  are  added  to 
diminish  resistance  and  to  improve  the  corrosion  of  the 
anode.  Since  the  solvent  action  of  the  acid  is  added  to 
the  corrosion  due  to  the  current,  acid  solutions  frequently 
show  an  efficiency  at  the  anode  greater  than  100  percent. 
By  dissolving  the  metal  already  deposited,  free  acid 
diminishes  the  current  efficiency  at  the  cathode.  With 
several  of  the  more  electro-positive  metals  having  a  low 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     89 

overvoltage  for  hydrogen,  it  is  probable  that  the  current 
efficiency  in  acid  solutions  is  also  diminished  by  the 
direct  deposition  of  hydrogen  instead  of  metal  by  the 
current. 

Addition  Agents 

The  use  of  addition  agents  has  already  been  referred  to 
page  85.  Rapid  progress  in  this  important  field  of 
study  is  now  being  made,  and  the  outlook  is  encouraging 
for  still  greater  improvement  in  the  future.  Some  of 
the  more  important  papers  upon  this  subject  follow: 

1.  Addition  agents  in  deposition  of  copper,  lead,  and  silver. 
Jarvis  and  Kern,  School  of  Mines  Quart.,  30:100-129. 

2.  Addition  Agents  in  Copper  Electrolytes  containing  Arsenic. 
Wen  and  Kern,  Tr.  Am.  Electrochem.  Soc.,  20:120-176. 

3.  Effect   of   Addition    Substances  in  Lead   Plating   Baths. 
Mathers  and  Overman,  Tr.  A.  E.  S.,  21  : 313-35. 

4.  Solid    Deposits   of   Lead   from   Lead   Acetate.     Mathers, 
Trans.  A.  E.  S.,  24:315-29. 

5.  Addition   Agents  in  the   Deposition  of  Iron.     Watts  and 
Li,  Trans.  A.  E.  S.,  25. 

6.  Addition  Agents  in  the  Deposition  of  Zinc  from  Zinc  Sul- 
phate.    Watts  and  Shape,  Trans.  A.  E.  S.,  25. 

Other  references  are  Nos.  2,  12,  13  and  14,  page  71  of  this 
book. 

The  Preparation  of  Metals  for  Plating 

The  preparation  of  the  surface  to  be  plated  is  a  very 
important  and  expensive  part  of  plating  operations. 
The  brief  directions  which  follow  are  intended  only  for 
laboratory  use.  For  the  practice  of  plating  establish- 
ments, standard  books  on  plating  should  be  consulted. 

General  Order  of  Operations 

Remove  scale,  if  present,  by  grinding  or  pickling  in 
acid  for  fifteen  to  thirty  minutes,  with  occasional 


90     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

scrubbing  with  a  metal-wire  brush,  rinse,  dip  in  lye, 
rinse  and  polish.  Hang  on  sling  wire,  remove  grease 
in  boiling  lye  or  the  electric  cleaner,  rinse  thoroughly, 
and  hang  in  the  plating  tank  immediately.  In  special 
cases  a  momentary  dip  in  weak  acid  or  an  amalgamating 
solution  may  be  required  after  the  removal  of  grease. 
Great  care  must  be  taken  that  none  of  the  solutions 
used  in  cleaning  are  carried  into  the  plating  baths,  or 
that  any  of  one  plating  solution  gets  into  another. 
Pickle  for  Removing  Scale  from  Iron  or  Steel. 

Sulphuric  acid,  cone.  1  part  by  weight 

Water  10  par"ts  by  weight 

Solution  for  Iron  Rust. 

Citric  acid  80  g. 

Ammonia  to  make  faintly  alkaline 
Water  to  make  one  liter. 

Requires  five  to  ten  hours  to  remove  thick  rust  with 
the  cold  solution,  and  ten  to  twenty  minutes  if  the 
solution  is  boiling.  The  iron  is  not  attacked  by  the 
fresh  solution. 

Bright  Dip  for  Brass. 

Sulphuric  acid,  cone.  100  g.        100  c.c. 

Nitric  acid,  cone.  75  g.        100  c.c. 

Salt  1  g.  2  g. 

Dry  brass  objects  are  immersed  in  this  for  one  second 
and  instantly  plunged  into  cold  water.  All  water  must 
be  excluded  from  the  bath. 

Lye  or  Caustic  Solution. 

Lye  1  Ib.  per  gallon  or  120  g.  per  liter. 
To  be  used  boiling  hot  for  the  removal  of  grease  from 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     91 

polished  articles.  Adhering  rouge,  etc.,  may  be  removed 
by  brushing  with  a  brush  made  of  cotton  twine,  known 
as  a  potash  brush.  Five  to  fifteen  minutes  may  be  re- 
quired for  the  removal  of  grease.  Only  animal  or  vege- 
table fats  are  saponified.  Mineral  oils  must  be  dis- 
solved by  benzine,  etc.  The  complete  removal  of  grease 
is  indicated  by  water  flowing  from  the  metal  in  an  even 
sheet.  The  student  should  test  the  flow  of  water  from 
the  metal  before,  and  several  times  during  treatment  with 
lye.  Zinc,  aluminum,  lead,  and  all  the  alloys  of  the 
latter  are  attacked  by  lye,  hence  these  materials  must 
be  immersed  only  momentarily  The  use  of  the  electric 
cleaner  is  preferable  for  such  materials. 

As  a  substitute  for  lye,  1  Ib.  per  gallon  of  the  "  Mineral 
Cleaner"  may  be  used.  It  is  claimed  that  this  removes 
even  mineral  oils,  and  does  not  injure  the  hands  or 
clothing. 

The  Electric  Cleaner. 

References : 

C.  F.  Burgess  in  Electrical  World,  1898,  Vol.  32,  page  445. 

C.  H.  Proctor  in  Metal  Industry,  Oct.,  1905. 

W.  L.  Churchill  in  Metal  Industry,  Aug.,  1908,  page  256. 

It  has  been  found  that  the  removal  of  grease  and  dirt 
from  metal  surfaces  by  hot  alkalies  is  greatly  facilitated 
by  using  the  object  as  anode  or  cathode  at  very  high 
current  densities  (40  to  80  amperes  per  square  foot) .  The 
action  seems  to  depend  mainly  on  the  tearing  off  of  the 
grease  by  the  storm  of  gas  bubbles  liberated  on  the 
surface  of  the  metal.  When  the  object  is  cathode,  the 
alkalinity  of  the  film  of  solution  in  contact  with  the  metal 
must  be  greatly  increased  by  electrolysis.  Some  prefer 
to  use  the  object  as  anode,  others,  as  cathode;  the 
majority  of  platers  favor  the  latter  practice.  With 
cathode  cleaning,  it  has  been  found  possible  to  remove 


92     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

grease  in  neutral  solutions  such  as  sodium  sulphate  or 
sodium  chloride.  This  may  be  advantageous  in  certain 
cases,  but  for  metals  which  are  not  attacked  by  lye,  it 
is  preferable  to  use  this,  or  the  "Mineral  Cleaner"  of 
half  the  strength  that  is  employed  without  the  current. 
The  author  has  found  the  latter  material  to  be  very 
satisfactory,  and  uses  it  to  the  exclusion  of  lye.  Small 
additions  of  potassium  cyanide  were  usually  made  to 
the  earlier  electric  cleaning  solutions  used  by  platers, 
but  this  dangerous  practice  has  now  become  less  common. 

Polishing. — Rough  surfaces  may  be  first  ground  on 
an  emery  wheel,  then  on  leather  or  felt  wheels  set  up 
with  fine  emery  powder  and  glue,  and  'finally  polished 
on  cloth  wheels  or  buffs,  touched  from  time  to  time  to 
a  block  of  tripoli  or  other  polishing  composition.  For  the 
highest  polish  after  plating,  rouge  may  be  used,  or  lamp 
black  made  to  a  paste  with  kerosene  or  alcohol.  De- 
posits of  silver,  gold  and  platinum  are  usually  burnished 
instead  of  buffed,  i.e.  rubbed  down  with  a  polished  tool 
of  steel  or  blood  stone,  lubricated  with  soap  suds.  This 
saves  the  usual  loss  of  metal  incurred  in  polishing,  and 
hardens  the  surface. 

In  plating  on  a  soft  metal  like  copper  with  nickel, 
cobalt  or  iron,  which  are  very  hard  when  deposited 
electroiytically,  much  labor  is  saved  by  giving  a  good 
polish  to  the  surface  to  be  plated. 

For  producing  a  bright,  metallic  luster,  either  before  or 
after  plating,  upon  objects  with  a  rough  surface,  wire 
wheels  or  hand  brushes  should  be  used.  Cloth  wheels 
polish  only  the  high  spots. 

The  remarks  which  follow  upon  plating  with  particular 
metals  call  attention  to  some  of  the  more  important 
points  to  be  observed,  but  for  complete  instructions, 
the  books  devoted  exclusively  to  plating  should  be 
consulted. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     93 

Nickel  Plating 
References : 

Electro-deposition  of  Nickel — Bennett,  Trans.  A.  E.  S., 

Vol.  25. 
Reactions  in  Electroplating — Brochet,   Metal  Industry, 

Oct.,  1908,  p.  314. 
Plating   with   Single   Salts— Brass    World,    June,    1907, 

page  196. 
Passivity    of    Nickel    Anodes — Schoch,  Amer.     Chem. 

Journal,  1909,  43,  235-56. 
Iron  in  the  Nickel  Deposit — Bancroft,  Trans.  A.  E.  S., 

Vol.  9,  page  217. 
Iron  in  the    Nickel   Deposit — Calhane   and    Gammage, 

Journal  Amer.  Chem.  S.,  1907,  29,  1268-74. 
Deposition  of  Nickel  on  Nickel — Snowdon,  Trans.  A.  E. 

S.,  7,  302. 
Deposition  of  Nickel  on  Nickel — Blasset,  Metal  Industry, 

Sept.,  1912,  page  375. 

For  other  references  on  nickeling,  see  the  references  on 
plating  in  general,  page  71. 

The  electro-deposition  of  nickel  presents  several  pecu- 
liarities. The  plating  bath  must  be  neutral,  or  but  very 
slightly  acid,  else  hydrogen  is  deposited.  The  reason 
for  this  is  apparent  when  the  discharge  potentials  of 
nickel  and  hydrogen  are  compared.  From  Caspari's 
value  for  the  overvoltage  of  hydrogen  on  nickel,  0.21, 
the  discharge  potential  of  hydrogen  on  nickel  would  be 
—  0.12  volts  (normal  calomel  electrode  =  —  0.56  volts). 
Some  values  for  the  discharge  potential  of  nickel  by 
different  observers3  are  0.75,  0.80,  0.903  volts.  Since 
the  discharge  potential  of  hydrogen  on  nickel  in  normal 
sulphuric  acid  is  much  lower  than  the  discharge  potential 
of  nickel  from  normal  nickel  sulphate,  it  is  evident 
that  if  much  sulphuric  acid  be  present,  the  deposit  should 
be  mainly  hydrogen.  The  necessity  of  excluding  from 
3  Electromotive  Force  of  Nickel — Schoch,  Amer.  Chem. 
Journal,  1909,  41,  208-31. 


94     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

the  nickel  bath  any  considerable  amount  of  a  strongly 
dissociated  acid  is  apparent. 

The  simultaneous  deposition  of  hydrogen  with  the 
nickel  causes  other  troubles  besides  a  low-current  effi- 
ciency. Pitting  is  troublesome  in  making  heavy  deposits 
of  nickel.  Here  and  there  bubbles  of  hydrogen  slowly 
form  on  the  cathode  and  cling  to  it,  cutting  off  the  cur- 
rent from  beneath  them,  and  as  the  surrounding  metal 
increases  in  thickness,  pits  are  produced. 

Surfaces  to  be  nickeled  must  be  moie  carefully  cleaned 
than  for  the  deposition  of  any  other  metal.  The  least 
trace  of  grease  or  tarnish  on  the  surface  may  cause  the 
plate  to  peel  off  and  roll  up  into  little  curls.  The  brittle- 
ness  and  tendency  to  curl  of  electrolytic  nickel  is  gen- 
erally and  probably  correctly,  ascribed  to  the  alloying 
of  hydrogen  with  the  metal. 

The  passivity  of  nickel  anodes  in  certain  solutions 
often  causes  trouble.  Rolled  (sheet)  anodes  are  es- 
pecially liable  to  become  passive  in  sulphate  solutions, 
resulting  in  the  production  of  free  acid,  which  in  turn 
affects  deposition  at  the  cathode.  Since  passivity 
increases  with  rise  of  current  density,  one  remedy  is 
to  employ  a  very  great  anode  surface.  It  has  been  found 
that  the  addition  of  a  little  chloride  to  the  bath  prevents 
passivity.  The  latter  remedy  according  to  Langbein, 
introduces  another  trouble  when  nickeling  iron.  It 
causes  rusting  of  the  iron  and  separation  of  the  nickel 
plate.  The  anodes  used  should  be  of  the  highest  purity 
attainable,  and  not  the  cast  anodes  usually  sold  by  dealers 
in  platers  supplies,  which  contain  only  ninety  or  ninety- 
two  percent  of  nickel,  the  remainder  consisting  of  iron, 
carbon,  and  copper  or  tin.  It  has  been  pointed  out 
that  too  small  an  anode  surface  causes  the  bath  to  become 
acid;  too  large  a  surface  may,  in  the  presence  of  the 
ammonium  salts  usually  used  in  nickel  baths,  give  it  an 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     95 

alkaline  reaction,  and  cause  the  production  of  a   dark 
and  poor  deposit. 

It  should  be  remembered  that  an  absolutely  clean 
surface  is  indispensable  for  nickel  plating,  and  once 
the  object  has  been  cleaned,  it  should  be  hung  in  the 
plating  bath  as  soon  as  possible  and  without  touching 
it  with  the  hands.  When  nickeling  brass  or  copper, 
the  last  thing  before  hanging  the  object  in  the  plating 
tank,  it  is  dipped  in  a  solution  of  potassium  cyanide  to 
dissolve  any  trace  of  oxide,  and  thoroughly  rinsed.  Iron 
is  often  coppered  before  nickeling,  but  it  can  be  nickeled 
directly  with  very  satisfactory  results  provided  it  is 
perfectly  clean. 

Copper  Plating 

Reference : 

The    Electro-deposition    of    Copper — C.    W.    Bennett, 
Trans.  A.  E.  S.,  Vol.  23,  pages  233-50. 

Baths  used  for  the  deposition  of  copper  are  of  two  kinds, 
acid  and  alkaline. 

The  Acid  Copper  Bath. — The  acid  copper  bath  is 
used  for  refining,  electrotyping,  and  wherever  a  thick 
deposit  or  rapid  deposition  is  desired.  It  cannot  be 
used,  however,  for  plating  directly  on  iron,  zinc  or  tin. 
The  student  should  consult  the  table  of  potentials  in 
sulphate  solutions  and  explain  the  failure  to  obtain  a 
satisfactory  deposit  on  these  metals.  It  was  found  that 
any  considerable  amount  of  free  mineral  acid  could  not 
be  used  in  the  nickel  bath.  The  amount  of  free  sul- 
phuric acid  used  in  copper  baths  varies  from  two  percent 
in  plating  solutions,  to  sixteen  percent  for  some  refining 
solutions.  The  use  of  this  large  amount  of  acid  is 
possible  because  copper  is  below  hydrogen  in  potential, 
and  is  but  slowly  attacked  by  sulphuric  acid.  Its 
purpose  is  to  diminish  resistance  and  improve  anode 
corrosion. 


96     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

The  Alkaline  Copper  Bath. — Alkaline  baths  are  of  two 
classes,  those  made  up  with  cyanide,  and  those  con- 
taining an  alkali  tartrate.  Since  the  former  are  used 
almost  exclusively,  the  others  will  not  be  considered  in 
commenting  on  the  alkaline  copper  bath. 

The  basis  of  the  bath  is  potassium  cyanide.  The 
cyanides  of  most  heavy  metals  are  insoluble  in  water, 
but  unite  with  potassium  cyanide  to  form  double  salts 
which  are  soluble.  The  copper  bath  contains  KCu 
(CN)2.  The  particular  advantage  of  the  cyanide  over 
the  acid  electrolyte  from  the  standpoint  of  the  plater, 
is  the  ability  to  plate  upon  iron,  zinc,  tin,  etc.  Another 
advantage  is  that  the  valence  of  copper  is  one  instead 
of  two,  and  therefore  twice  as  much  copper  should  be 
deposited  by  the  same  current  as  in  the  acid  bath. 
The  continuous  evolution  of  hydrogen  cuts  down  the 
current  efficiency,  however,  and  when  heavy  deposits 
are  desired,  after  the  object  has  been  plated  for  fifteen  to 
thirty  minutes  in  the  alkaline  solution,  it  is  removed  to 
the  acid  bath.  Explain  the  evolution  of  hydrogen  in  the 
alkaline  bath,  and  its  failure  to  appear  in  the  acid 
solution.  The  copper  cyanide  solution  plates  into 
hollows  and  cavities  much  better  than  the  acid  bath, 
or,  to  use  the  technical  term,  " throws"  better. 

The  chemical  action  at  the  anode  consists  in  the 
union  of  the  cyanogen  liberated  by  the  current  with  the 
anode,  forming  cuprous  cyanide.  Since  this  is  insoluble 
in  water,  it  forms  a  coating  over  the  anode  which  would 
finally  stop  the  current  if  not  removed.  It  is  therefore 
necessary  to  have  potassium  cyanide  in  contact  with  the 
anode  to  combine  with  the  cuprous  cyanide  and  convert 
it  into  soluble  KCu(CN)2.  Cyanide  is  therefore  added 
in  excess  of  that  required  to  form  the  double  cyanide 
with  all  the  copper  originally  in  the  bath.  This  is 
referred  to  as  "free  cyanide."  Since  a  solution  of 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     97 

cyanide  is  slowly  converted  to  potassium  carbonate 
and  other  compounds  by  the  carbon  dioxide  of  the 
air,  it  is  necessary  to  add  a  little  cyanide  occasionally. 

A  trouble  to  which  cyanide  baths  are  subject  is  " spot- 
ing  out."  Some  weeks  or  even  months  after  plating, 
small  tarnished  spots  appear,  injuring  the  appearance, 
and  in  the  case  of  silvered  mirrors,  the  usefulness  also. 
In  the  case  of  plating  on  cast  iron,  at  least,  the  trouble 
seems  to  be  due  to  electrolyte  contained  in  minute  cavities 
in  the  metal.  The  remedy  is  to  destroy  the  cyanide  by 
repeatedly  dipping  in  dilute  acetic  acid,  alternated  with 
washing  in  hot  water.  In  other  cases,  spotting  out  would 
seem  to  be  due  to  the  settling  of  the  chemical  dust  of  the 
plating  room  upon  the  work  after  polishing. 

The  student  should  remember  that  cyanide  solutions 
are  extremely  poisonous.  Objects  taken  from  cyanide 
baths  should  not  be  handled  until  thoroughly  rinsed. 
The  hands  should  be  thoroughly  washed  before  leaving 
the  laboratory  after  working  with  cyanide  plating 
baths. 

The  Deposition  of  Alloys 

References : 

Conditions  which  Determine  the  Composition  of  Electro- 
deposited  Alloys — S.  Field,  Trans.  Faraday  Soc.,  Sept., 
1909,  Vol.  5,  pages  172-94. 

Electrolytic  Precipitation  of  Bronzes — Curry.,  Trans. 
A.  E.  S.,  Vol.  9,  pages.  249-53. 

Electrolytic  Deposition  of  Zinc-Nickel  Alloys — Schoch, 
Trans.  A.E.S.,  Vol.  11,  pages  135-51. 

Watt's  Electro-Deposition,  pages  534—36. 

From  experiments  with  a  cyanide  brass  bath  referred 
to  above,  Field  says,  "It  is  then  found  that,  with  a 
solution  containing  about  equal  quantities  of  the  two 
salts  in  the  absence  of  any  notable  amount  of  free 
cyanide. 


98     A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

(1)  Copper  is  more  readily  deposited,  and 

(2)  The  percentage  of  zinc  increases  with  the  current 
density. 

(3)  The  percentage  of  zinc  increases  as  the  amount  of 

zinc  compound  is  increased,  while 

(4)  Even  with  a  large  excess  of  zinc  compound  deposits 

containing  a  fair  proportion  of  copper  are  readily 
obtained. 

(5)  The  effect  of  dilution  is  to  raise  the  percentage  of 

zinc,  on  account  of  the  higher  E.M.F.  necessary 
to  maintain  the  same  current  density,  while 

(6)  A  rise  of  temperature  induces  the  deposition  of  a 

larger  proportion  of  copper. 

(7)  With   appreciable   amounts   of  free   cyanide   the 

percentage  of  copper  is  always  high,  even  with 
high  current  density,  while  the  free  cyanide  adds 
little  conductance  to  the  solution,  other  than 
that  it  prevents  the  formation  of  insoluble  single 
cyanides  at  the  anodes." 

He    summarizes    his    experiments    as  follows:     ''The 
conclusions  drawn  from  these  results  are — 

1.  Brass   is   deposited   quantitatively,    or   nearly    so, 
over  wide   ranges   of   composition  from   a  mixture   of 
cyanides. 

2.  Even  in  cyanide  solutions  the  greater  ease  of  deposi- 
tion of  copper  is  marked  in  the  absence  or  excess  of 
free  cyanide. 

3.  Conditions  tending  to  raise  the  E.M.F.  increases 
the  percentage  of  zinc,  such  as — 

a.  Dilution  of  solution 

b.  Decrease  of  temperature. 

4.  Anodes    are   freely    soluble    with    warm,    agitated 
solution  even  in  the  presence  of  only  small  amounts  of 
free  cyanides. 

5.  The  effect  of  free  cyanide  is  to 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY     99 

a.  Make  zinc  more  positive  with  respect  to  copper, 
and  thereby  increase  the  percentage  of  copper 
in  deposits. 

b.  Increase  the  evolution  of  hydrogen,  and 

c.  Induce  abnormal  anode  efficiencies." 

The  author  cannot  subscribe  to  Field's  third  conclusion. 
The  E  M.F.  per  se,  has  nothing  whatever  to  do  with  the 
composition  of  the  deposit. 

From  theoretical  considerations,  the  following  con- 
ditions appear  to  be  of  prime  importance  in  the  deposi- 
tion of  alloys  of  two  metals : 

1.  The  potentials  of  the  metals  in  the  particular  electro- 
lyte chosen  should  be  near  together  in  order  that  an  alloy 
may  be  successfully  deposited. 

2.  By   regulation   of   the   relative   concentrations    of 
the  metals,  and  of  the  current  density,  the  composition  of 
the  alloy  may  be  controlled.     Dilution  of  the  solution  or 
increase  of  current  density  will  cause  the  deposit  to  contain 
a  larger  proportion  of  the  more  electro-positive  metal. 

3.  Stirring  causes    the  deposition  of    a  smaller  pro- 
portion of  the  more  electro-positive  metal  (zinc),  and  has 
the  same  effect  on  the  composition  as  increasing  the  con- 
centration or  lowering  the  current  density. 

4.  Increase  of  temperature  in  stationary  electrolytes 
increases  mobility,  and  therefore  should  have  the  same 
effect  as  stirring — provided  the  composition  of  the  electro- 
lyte is  not  changed  at  the  same  time  through  increased 
anode  corrosion.    Pfanhauser,  however  (page  353),  states 
that  the  opposite  occurs  in  the  brass  bath. 

5.  A  moderately  dilute  electrolyte  is  a  necessity  under 
present  current  conditions.     In  the  deposition  of  a  single 
metal,  a  high  content  of  metal  in  the  electrolyte  is  desir- 
able since  it  permits  of  more  rapid  deposition  than  from  a 
dilute  solution.     A  case  in  point  is  the  modern  "  hi-power" 
nickel  bath.     In  depositing  alloys  the  use  of  very  con- 


100  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

centrated  electrolytes  will  result  in  an  alloy  containing 
too  large  a  proportion  of  the  electro-negative  metal.  In- 
crease of  the  current  density  for  the  purpose  of  securing 
the  desired  proportion  of  the  electro-positive  metal  in  the 
alloy  will  probably  result  in  a  " burned"  deposit  before 
the  proper  composition  is  attained.  This  should  not  be 
taken  to  mean  that  no  increase  in  concentration  over 
present  practice  is  possible;  but  that  such  increase  is  more 
limited  than  in  the  deposition  of  a  single  metal,  and  with 
the  stronger  solutions,  a  higher  current  density  must  al- 
ways be  used.  The  farther  apart  the  potentials  of  the  two 
metals,  the  more  dilute  must  be  the  electrolyte. 

6.  For  proper  control  of  the  composition!  of  the  alloy  in 
practical  plating  it  is  desirable  that  the  metals  and  alloy 
shall  differ  in  color.  The  failure  of  several  processes  for 
the  electro-deposition  of  alloys,  which  otherwise  seemed  to 
have  good  prospects  for  commercial  success,  may  be  laid 
to  their  not  fulfilling  the  above  condition. 

Of  the  many  alloys  which  have  been  deposited  electro- 
lytically,  brass  alone  is  of  commercial  importance. 

Brass  Plating 

The  brass  bath  has  the  reputation  of  being  the  most 
difficult  of  commercial  baths  to  control  so  that  it  will 
yield  good  deposits  of  uniform  and  satisfactory  color. 
The  ratio  of  copper  to  zinc  varies  widely  in  solutions 
recommended  by  different  platers,  but  equal  weights 
gives  a  bath  that  is  probably  as  easy  to  control  as  any. 

Anodes  should  equal  or  exceed  the  cathodes  in  surface, 
and  should  be  of  cast  brass.  On  account  of  the  unsatis- 
factory corrosion  of  brass  anodes,  some  platers  re  commend 
the  use  of  copper  anodes  and  the  addition  of  zinc  salts 
from  time  to  time.  Brass  anodes  are  always  left  in  the 
bath. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  101 

Too  little  free  cyanide  causes  a  reddish  tone  to  the 
deposit,  and  the  formation  of  a  thick  coating  on  the 
anodes. 

Too  much  cyanide  causes  a  vigorous  evolution  of  gas  at 
the  cathode,  and  a  slow  forming  deposit  which  peels  under 
the  scratch  brush.  Excess  of  cyanide  is  particularly  to  be 
avoided  for  plating  iron  objects. 

Langbein  advises  adding  115  g.  per  liter  of  sodium 
carbonate  for  maintaining  a  bright  brass  color  when  plat- 
ing iron  or  zinc.  Old  baths  are  likely  to  contain  much 
carbonate. 

Warm  baths  (40°-50°  C.)  are  claimed  to  give  much 
better  results  than  cold  solutions.  Rise  of  temperature 
should  help  both  circulation  and  anode  corrosion. 

Silver  Plating 

Except  for  iron,  antimony  and  lead,  the  metals  which  it 
is  desirable  to  plate  with  silver  are  electro-positive  to  this 
metal  in  cyanide  solutions.  The  result  is  that  special 
methods  must  be  used  in  plating  these  metals  with  silver 
else  the  bath  is  decomposed  and  the  object  receives  by 
immersion  a  poorly  adherent  coating,  which  later  causes 
the  whole  deposit  to  come  off. 

One  of  the  processes  used  to  secure  better  adhesion  is 
"quicking."  Of  this  Field  (page  228)  says: 

"Quicking — this  term  denotes  a  further  preliminary 
operation  of  passing  the  prepared  work  through  a  solu- 
tion of  mercury  cyanide.  Mercury  is  less  positive  than 
either  copper  or  silver.  In  such  a  solution,  copper, 
brass  and  German  silver  receive  a  bright  deposit  of  mer- 
cury, and  when  thus  covered,  the  work  is  thoroughly 
rinsed,  and  transferred  to  the  silver  bath,  when,  on  account 
of  mercury  being  less  positive  than  silver,  no  deposition 
of  silver  by  simple  immersion  can  occur,  and  the  sub- 


102  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

sequent  electrolytic  deposit  is  more  adherent."  Others 
ascribe  the  beneficial  effect  of  "quicking"  entirely  to  an 
alloying  of  the  mercury  with  both  metals.  Thin  objects 
should  be  immersed  in  the  quick  dip  only  a  few  seconds  as 
a  heavy  deposit  of  mercury  may  make  them  brittle.  The 
quicking  solution  must  be  entirely  rinsed  from  the  objects 
before  transferring  them  to  the  silver  bath,  and  if  the 
amalgamated  article  is  a  dull  gray,  it  should  be  rubbed 
bright  with  a  brush.  The  quick  dip  is  used  on  copper, 
brass  and  German  silver,  but  is  said  to  be  useless  for  iron, 
steel  or  nickel  objects. 

An  alternative  process  to  " quicking"  is  the  use  of  a 
special  ''striking  bath"  (No.  17,  page  77)  to  form  by  the 
current  a  thin  film  of  silver  before  plating  in  the  regular 
bath.  The  striking  bath  is  very  weak  in  silver  and  contains 
a  large  excess  of  cyanide.  The  object  is  plated  in  this 
for  only  a  few  seconds. 

Langbein  says  that  copper,  brass,  bronze  and  German 
silver  may  be  silvered  directly,  but  that  iron,  steel, 
nickel,  zinc,  tin,  lead  and  Britannia  should  be  coppered 
or  brassed  first.  Tin  and  Britannia  may  be  directly 
silvered  by  the  use  of  a  striking  bath. 

Silver  Direct  on  Steel  vs.  Preliminary  Nickeling — 
"All  large  manufacturers  of  silver-plated  steel  cutlery 
deposit  silver  directly  on  steel.  Several  smaller  firms 
nickel  first.  Evidence  is  in  favor  of  the  former  practice, 
as  many  of  the  large  firms  used  to  nickel."4 

It  is  recommended  that  the  anodes  of  purest  sheet 
silver  be  hung  by  iron  wire.  The  current  should  be  cut 
off  before  removing  objects  from  the  bath,  else  the  de- 
posit may  have  a  yellowish  tone. 

Compare  the  evolution  of  hydrogen  with  that  in  the 
other  cyanide  baths  (copper  and  brass)  and  explain. 

4  Brass  World,  1912,  page  135. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  103 

EXPERIMENTS  IN  PLATING 

Deposition  by  Immersion 

EXPERIMENT  66 

DEPOSITION  OF  COPPER  ON  VARIOUS  METALS  BY 
IMMERSION 

Prepare  500  c.c.  of  a  nearly  saturated  solution  of  copper 
sulphate.  Clean  small  sheets  of  aluminum,  brass,  copper, 
iron,  lead,  nickel,  tin  and  zinc.  Dip  each  for  a  few 
seconds  in  the  solution,  examine  and  test  the  adhesion  of 
the  deposit,  if  any.  Explain. 

EXPERIMENT  67 

DEPOSITION    BY    IMMERSION    FROM    THE    LABORATORY 
PLATING  BATHS 

Test  the  action  of  the  metals  of  the  foregoing  experi- 
ment on  the  laboratory  plating  baths  for  the  deposition  of 
brass,  copper  (acid  and  alkaline),  lead,  nickel  and  zinc. 

EXPERIMENT  68 

THE  COPPERING  OF  IRON  BY  IMMERSION 

Prepare  a  solution  containing  50  g.  of  copper  sul- 
phate and  50  g.  of  concentrated  sulphuric  acid  per 
liter.  This  has  been  recommended5  for  giving  iron 
a  light  coating  of  copper.  Test  it  by  immersing  clean 
wire  nails.  What  effect  has  the  time  of  immersion? 
How  does  this  deposit  compare  with  that  from  a  satur- 
ated solution  of  copper  sulphate  in  appearance?  In 
adhesion?  Explain. 

Gore's  experiments  on  deposition  by  immersion  are 
given  in  table  7. 

5  Langbein,  6th  edition,  page  487. 

8 


104  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

Plating  by  Contact. — This  method  was  formerly  much 
used  by  platers.  It  consists  in  placing  in  the  plating  solu- 
tion, in  contact  with  the  metal  to  be  plated,  pieces  of 
some  more  electro-positive  metal,  generally  aluminum 
or  zinc.  The  bath  is  usually  heated. 

EXPERIMENT  69 
NICKELING  BY  CONTACT 

Heat  to  70°  or  80°  C.  a  solution  containing  per  liter 
30  g.  nickel  sulphate  and  30  g.  ammonium  chloride. 
In  this  immerse  the  cleaned  articles  to  be  plated,  in 
contact  with  several  pieces  of  aluminumi  Stir  occasion- 
ally. After  five  minutes,  remove  the  articles  and  test  the 
durability  of  the  deposit  on  the  polishing  wheel.  Try 
using  zinc  instead  of  aluminum. 

EXPERIMENT  70 
TINNING  BY  CONTACT 

This  process  is  used  for  the  tinning  of  brass  pins  and 
other  small  objects.  Boil  several  small  pieces  of  brass 
or  copper  wire  in  contact  with  granulated  tin  in  a  solu- 
tion of  cream  of  tartar  to  which  a  little  stannous  chloride 
has  been  added.  For  a  thick  deposit,  zinc  should  be 
used  in  place  of  the  granulated  tin.  Why?  Langbein, 
page  502,  gives  as  a  solution  for  tinning  by  zinc  contact 
10  g.  cream  of  tartar  and  27  g.  of  stannous  chloride  per 
liter. 

State  the  advantages  of  plating  by  contact.  Its 
disadvantages. 

Under  the  names  "Galvanite,"  "Voltite"  and  "  Nick- 
elite,  "  several  plating  powders  have  been  put  upon  the 
market,  which  for  their  operation  depend  on  plating  by 
contact. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  105 

Galvanite. — (Electrical  Review  1910,  Vol.  66  page 
243,  and  Brass  World,  1910,  page  124,  177,  365.) 

This  is  a  paste  consisting  of  the  metal  (or  one  of  its 
salts)  which  is  to  be  deposited,  an  electro-positive  metal 
such  as  aluminum,  zinc,  cadmium  or  magnesium  in  a 
very  finely  divided  state,  a  substance  capable  of  produc- 
ing an  aqueous  electrolyte  when  brought  into  contact 
with  moisture,  and  inert  substances  such  as  chalk,  soap- 
stone,  kieselguhr,  boric  acid,  dextrine,  etc.,  to  prevent  too 
rapid  action,  and  to  act  as  polishing  powders.  The 
article  to  be  plated  is  rubbed  with  a  damp  cloth  containing 
the  paste. 

GALVANITE  FOR  NICKELING 

Nickel  ammonium  sulphate  60  percent 

Magnesium  3  percent 

Chalk  30  percent 

Talc  powder  7  percent 

FOR  GALVANIZING 

Ammonium  sulphate  15  percent 

Zinc  dust  45  percent 

Magnesium  3  percent 

Chalk  30  percent 

Talc  powder  7  percent 

Voltite. — The  results  of  gold  and  silver  plating  with 
this  are  described  in  The  Metal  Industry,  1912,  pages 
77,  261,  and  it  is  claimed  that  application  of  voltite  for 
five  minutes  gave  a  deposit  of  silver  which  it  would  re- 
quire four  or  five  hours  to  deposit  in  the  regular  silver 
plating  bath. 

EXPERIMENT  71 

GALVANITE  NICKELING  BY  ZINC  DUST 

Prepare  galvanite  nickeling  paste,  using  zinc  dust  as 
the  electro-positive  metal,  and  plate  pieces  of  polished 


106  A  LABORATOKY  COURSE  IN  ELECTROCHEMISTRY 

brass  and  copper.  Note  the  time  required  for  a  good 
looking  coating,  and  plate  similar  pieces  for  the  same 
time  in  the  regular  nickel  bath.  Test  the  deposits  with 
the  polishing  wheel. 

EXPERIMENT  72 

GALVANITE  NICKELING  BY  PULVERIZED  ALUMINUM 

Prepare  galvanite  nickeling  paste  with  aluminum 
bronze  as  the  electro-positive  metal.  Use  and  test  as 
in  experiment  71.  Equal  parts  of  nickel  sulphate 
and  ammonium  chloride  will  work  better  than  nickel 
ammonium  sulphate.  Why? 

EXPERIMENT  73 
GALVANITE  TINNING 

Prepare  and  test  a  galvanite  tin  paste  after  formula 
No.  20,  page  78,  with  zinc  dust.  It  may  be  necessary  to 
add  some  salt  to  accelerate  corrosion  of  the  zinc;  what 
would  you  suggest? 

Plating  by  Electro -deposition 

In  experiments  on  plating,  each  student  should  re- 
cord data  such  that  another  could  repeat  the  experi- 
ment and  obtain  the  same  result.  This  means  securing 
the  facts  in  regard  to  the  important  factors  enumerated 
on  page  84. 

EXPERIMENT  74 

THE   EFFECT   OF   CURRENT   DENSITY  IN   ELECTRO- 
DEPOSITION 

.Test  the  effect  of  current  density  in  the  deposition  of 
nickel  from  the  standard  bath,  No.  8.  Cut  a  sheet  of 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY   107 

copper  or  brass  to  fill  one  end  of  a  rectangular  glass  jar, 
polish,  and  with  a  knife  mark  a  series  of  horizontal 
lines  across  it,  2  cm.  apart.  Clean  the  sheet  and 
immerse  it  to  one  of  the  upper  marks.  It  should  fit 
snugly  against  the  glass,  so  that  the  current  is  confined 
to  one  side.  Immerse  the  anode  to  the  same  depth  at 
the  other  end  of  the  jar.  For  five  minutes  plate  at  the 
current  density  specified  for  nickeling.  Examine.  Raise 
the  cathode  to  the  next  mark,  raise  the  anode  also,  and 
increase  the  current  fifty  percent.  Continue  in  this 
way  until  the  bottom  of  the  cathode  is  reached.  Com- 
pute the  current  densities,  and  find  that  at  which  the 
deposit  first  became  dark,  Each  deposit  should  be 
examined  before  proceeding  with  the  next.  Where  does 
the  first  dark  deposit  appear?  Why  there? 

EXPERIMENT  75 

CIRCULATION  AND  CRITICAL  CURRENT  DENSITY 

Test  the  effect  of  stirring  on  the  solution  just  used, 
by  putting  a  mechanical  stirrer  into  the  cell  and  repeat- 
ing the  experiment. 

EXPERIMENT  76 

CONCENTRATION  AND  CRITICAL  CURRENT  DENSITY 

Determine  similarly  the  critical  current  density  for 
the  "hi-power"  Prometheus  nickel  solution,  without 
stirring. 

EXPERIMENT  77 
CRITICAL  CURRENT  DENSITY  IN  THE  BRASS  BATH 

After  the  manner  of  experiment  74,  determine  the 
critical  current  density  for  the  brass  bath.  Note  the 
effect  of  current  density  on  the  color  of  the  deposit. 


108  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

EXPERIMENT  78 

THE  EFFECT  OF  CIRCULATION  ON  THE  COLOR  OF  BRASS 
PLATING 

Test  the  effect  of  stirring  on  the  color  of  the  brass 
deposit.  Connect  two  jars  of  the  brass  bath  in  series, 
put  a  mechanical  stirrer  near  the  cathode  in  one  cell, 
and  plate  at  0.3  amperes  per  sq.  dm.,  or  whatever  may 
be  the  least  current  that  gives  a  good  yellow  brass  in 
the  stationary  solution.  Compare  the  deposits.  Repeat 
at  double  the  former  current.  At  four  times  the  initial 

current.     Explain  the  results. 

* 

EXPERIMENT  79 
NICKELING  POINTED  OBJECTS 

Compare  the  suitability  of  nickel  bath  No.  8,  or  No.  9 
with  No.  10  for  plating  on  pointed  objects.  Cut  equal 
narrow  triangles  of  sheet  copper  or  brass,  and  plate  them 
on  both  sides  for  twenty  to  thirty  minutes  with  the  two 
baths  in  series.  How  can  you  arrange  the  anodes  to  assist 
in  securing  an  even  distribution  of  the  deposit?  At  the 
same  time  plate  a  similar  strip  in  the  large  nickel  tank  in 
the  plating  room,  without  taking  any  precautions  about 
arrangement  of  the  anodes.  The  current  density  should 
be  the  same  in  the  large  tank  as  in  the  others. 

EXPERIMENT  80 

THE  PARTICULAR  REQUIREMENT  IN  A  BATH  FOR 
POINTED  OBJECTS 

Why  is  bath  No.  10  especially  recommended  for  pointed 
objects?  Is  it  due  to  some  specific  action  of  the  citrate, 
or  merely  because  the  resistances  of  the  two  baths  are 
different?  The  practical  plater  controls  the  bath  by 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  109 

regulation  of  voltage,  and  rarely  knows  what  current 
he  is  using.  Measure  the  resistivity  of  each  bath, 
dilute  one  to  the  resistivity  of  the  other,  and  repeat 
79,  having  conditions  at  the  two  cathodes  as  exactly 
alike  as  possible.  Before  closing  the  switch,  consult 
the  laboratory  instructor  to  be  sure  that  you  have 

secured  the  condition  last  specified. 

• 

EXPERIMENT  81 
NICKEL    PLATING    ON    ZINC 

Test  baths  No.  8  or  No.  9  and  No.  10  for  the  direct 
nickeling  of  zinc.  The  baths  should  be  in  separate  cir- 
cuits instead  of  in  series  as  in  previous  cases.  Exercise 
your  ingenuity  to  secure  good  deposits  from  both  baths. 
What  is  the  cause  of  the  difficulty  in  nickeling  zinc? 
In  what  other  ways  might  it  be  remedied? 

EXPERIMENT  82 

THE  EFFECT  OF  TEMPERATURE  UPON  THE  QUALITY  OF 
ELECTRO-DEPOSITED  NICKEL 

Compare  the  physical  qualities  of  electrolytic  nickel 
deposited  at  room  temperature,  and  at  65°  to  70°  C. 
Use  bath  No.  9,  or  solution  from  the  large  plating  tank. 
Operate  the  two  cells  in  series  for  forty-five  minutes  or 
more  at  a  current  density  of  0.5  amperes,  setting  one  jar  in 
a  tin  box,  heated  by  electric  lights.  Test  the  deposits  for 
hardness,  brittleness  and  adhesion. 

EXPERIMENT  83 

UNEVEN  DISTRIBUTION  OF  CURRENT  OVER  THE  CATHODE 

Determine  the  distribution  of  current  over  a  flat 
cathode  by  measuring  the  thickness  of  a  piece  of  sheet 


110  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

copper  before  and  after  plating  in  the  acid  copper  bath. 
Use  the  large  plating  solution,  and  deposit  at  a  current 
density  of  1  ampere  for  sixty  to  ninety  minutes. 
Show  the  distribution  of  current  by  a  chart.  How  can  a 
more  even  distribution  of  current  be  secured  in  plating? 
Uneven  distribution  of  current  is  also  objectionable  in 
electrolytic  refining,  where  cathodes  remain  in  the  elec- 
trolyte for  many  days.  Can  you  suggest  precautions 
for  minimizing  this  trouble? 

EXPERIMENT  84 

• 
CONCENTRATION  CHANGES  DURING  ELECTROLYSIS 

Study  the  changes  in  concentration  which  occur  in  a 
plating  solution.  This  may  be  learned  by  placing  the 
solution  in  a  glass  cell  in  front  of  a  strong  light  and 
examining  it  during  passage  of  the  current.  After 
an  hour's  operation,  considerable  changes  may  have 
occurred.  Is  this  action  helpful  or  harmful  in  plating? 
In  refining? 

EXPERIMENT  85 
PLATING  IRON  WITH  COPPER 
Clean  a  piece  of  iron  and  plate  it  with  copper. 

PLATING  ON  ALUMINUM 

References : 

Langbein,  6th  Ed.,  page  471 

Wisconsin  Engineer,  Jan.,  1912,  pages  162-5. 

A.  Fischer,  Electrochemical  Industry,  1903,  page  584. 

Burgess    and    Hambuechen,    Electrochemical    Industry, 

1904,  page  85. 

A.  Lodyguine,  Trans.  A.  E.  S.,  7,  153. 
Loeb's  Method,  Brass  World,  1909,  page  145. 
Szarvasy's  Method,  Brass  World,  1909,  page  280. 
O.  Meyer,  Metallurgical  and  Chem.  Eng.,  1908,  6,  510. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  111 

EXPERIMENT  86 
ELECTROPLATING  ALUMINUM 

Clean  a  sheet  of  aluminum  4  by  15  cm.  in  the  electric 
cleaner,  and  plate  it  in  the  acid  copper  bath  for  twenty 
minutes.  Plate  another  piece  of  aluminum  in  the  alka- 
line copper  bath.  Cut  away  all  the  edges  and  try  strip- 
ping the  deposits.  Examine  for  porosity  by  transmitted 
light.  Try  again,  dipping  the  aluminum  in  dilute  sul- 
phuric acid  and  rinsing,  after  treatment  in  the  electric 
cleaner.  Try  dilute  hydrofluoric  acid  similarly.  In 
which  case  is  the  adhesion  best? 

EXPERIMENT  87 
NICKEL  PLATING  ON  ALUMINUM 

Plate  aluminum  with  nickel  after  the  method  of 
experiment  86. 

EXPERIMENT  88 

PLATING  ALUMINUM  WITH  IRON 
Plate  aluminum  with  iron  by  the  above  method. 

EXPERIMENT  89 
PLATING  ALUMINUM  WITH  ZINC 

Plate  aluminum  with  zinc  by  the  above  method. 
Compare  the  adhesion  and  porosity  of  the  deposits  in 
experiments  86-89.  Comment. 

EXPERIMENT  90 

THE    CURRENT    EFFICIENCY    OF    THE  DEPOSITION    OF 

COPPER 

Prepare  a  copper  coulombmeter6  by  placing  two  anodes 
in  a  glass  cell  containing  per  liter  150  g.  copper  sulphate, 

6  Oettel,  Exercises  in,  Electrochemistry,  pages  16  and  23. 


112  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

50  g.  sulphuric  acid,  and  50  g.  of  alcohol.  Connect  in 
series  the  coulombmeter,  a  cell  containing  the  copper 
cyanide  plating  bath,  an  accurate  ammeter  and  two  rheo- 
stats. Electrolyze  for  an  hour  at  a  current  density  of  0.5 
amperes,  keeping  the  current  constant  by  adjustment  of 
the  rheostats.  These  should  be  chosen,  one  for  coarse, 
the  other  for  fine  adjustment  of  the  current.  Use  two 
anodes,  and  a  carefully  weighed  cathode  in  each  cell. 
At  the  end  of  the  time,  open  the  switch,  quickly  remove 
the  cathodes,  wash,  rinse  in  distilled  water,  then  in  alco- 
hol, dry  by  hot  air,  and  weigh.  From  the  ampere-hours, 
compute  by  Faraday's  law  the  theoretical  amount  of 
deposit,  and  the  current  efficiency  in  eac*h  cell.  Oettel 
gives  1.182  g.  per  ampere-hour  for  the  coulombmeter, 
instead  of  1.186  g.  required  by  Faraday's  law.  How 
does  your  result  compare  with  this? 

EXPERIMENT  91 

THE  CURRENT  EFFICIENCY  OF  THE  NICKEL  BATH 

Using  the  copper  coulombmeter  and  ammeter,  find 
the  current  efficiency  of  deposition  of  the  nickel  bath  used 
in  the  laboratory. 

EXPERIMENT  92 

THE   CORROSION   OF   ALUMINUM   ELECTRODES 

With  the  coulombmeter  or  an  ammeter,  electrolyze 
in  series  for  an  hour  at  0.5  amperes  per  sq.  dm.  fifteen 
percent  solutions  of  sodium  chloride,  and  of  potassium 
sulphate  with  cleaned  and  carefully  weighed  aluminum 
electrodes,  the  former  at  room  temperature,  and  the 
latter  at  50°  to  60°  C.  Use  the  110-volt  circuit  and 
suitable  rheostats  for  accurate  control  of  the  current. 
The  electrodes  should  be  marked  for  identification  before 
cleaning.  Calculate  the  current  efficiency  at  each 
electrode.  Explain. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY   113 

(         EXPERIMENT  93 

THE  RELATIVE  PROTECTION  AFFORDED  BY  DEPOSITS  OF 

LEAD  AND  COPPER 

Plate  a  cleaned  sheet  of  iron  with  lead  from  the  silico- 
fluoride  solution  for  a  half  hour,  and  plate  similarly 
another  sheet  with  copper.  Bend  up  the  edges  to  form 
trays,  put  a  little  ten  percent  sulphuric  acid  in  each  and 
observe  the  protection  against  corrosion  of  the  iron 
afforded  by  the  two  coatings. 

EXPERIMENT  94 
THE  DEPOSITION  OF  BRASS 

a.  Insert  a  brass  anode  in  a  two  percent  solution  of 
copper  sulphate  (sp.  gr.  1.0126  at  18°  C.)  and  electrolyze 
long  enough  to  determine  the  character  of  the  deposit. 
Add  3  g.  of  zinc  sulphate  per  100  c.c.  and  repeat.     Con- 
tinue additions  of  zinc  sulphate  in  amounts  of  3  g.  up  to 
12  g.  per  100  c.c.     Record  the  current  and  current  density 
used  in  each  case,  varying  this  as  you  think  is  best  suited 
to  the  production  of  brass.     Your  conclusions  about  the 
deposition  of  brass  from  a  mixture  of  the  sulphates  of 
zinc  and  copper? 

b.  Now  dissolve  2  g.  copper  sulphate  and  2  g.  zinc 
sulphate  in  100  c.c.  of  water,  stir  in  5  g.  dry  sodium  car- 
bonate, and  slowly  stir  in  a  solution  of  10  g.  potassium 
cyanide  in  50  c.c.  water,  until  the  blue  color  of  the  copper 
salt  has  disappeared.     Test  this  for  the  deposition  of 
brass. 

Explain  the  results  in  (a)  and  (b). 

EXPERIMENT  95 

THE  DEPOSITION  OF  A  NICKEL-!RON  ALLOY,  ABOUT 

FIFTY  PERCENT 

First  consult  the  table  of  potentials  and  report  to  the 
instructor  on  the  most  promising  electrolyte  and  con- 


114  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

centration  to  use.  Then  test  the  electrolyte  at  various 
current  densities.  Does  the  deposit  contain  both  metals? 
How  do  the  deposits  compare  with  electrolytic  iron  and 
nickel  in  brittleness,  curling,  hardness,  rusting,  etc.? 

EXPERIMENT  96 
DEPOSITION    OF    A    CADMIUM-COPPER    ALLOY 

Make  a  preliminary  report  as  in  experiment  95.  In 
testing  the  electrolyte,  start  with  the  copper  solution  of 
known  strength,  and  to  it  add  the  weighed  cadmium  salt 
in  small  amounts,  testing  the  deposit  after  each  addition. 
How  does  the  process  compare  with  the  electro-deposi- 
tion of  the  series  of  zinc-copper  alloys? 

EXPERIMENT  97 

THE   DEPOSITION  OF  A  TIN-CADMIUM  ALLOY 

All  tin-cadmium  alloys  should  be  white,  and  the 
cadmium-rich  alloys  may  possibly  make  a  good  protec- 
tive coating  for  iron.  The  alloys  in  the  middle  of  the 
series  may  prove  much  harder  than  either  metal,  and 
will  probably  tarnish  less  readily. 

Make  a  preliminary  report  on  promising  electrolytes. 
If  any  can  be  found,  try  them. 

EXPERIMENT  98 

"  ARC  AS"  SILVER  PLATING — AN  ALLOY  OF  SILVER  AND 
CADMIUM.     WATT-PHILIP,  PAGE  473 

The  bath  is  patented  in  England  (No.  1391  of  1892) 
by  S.  O.  Cowper-Coles,  and  is  claimed  to  produce  an 
alloy  that  is  harder,  and  tarnishes  less  readily  than  pure 
silver.  To  make  the  bath,  for  each  liter,  dissolve  3.75  g. 
of  silver  and  86.7  g.  of  cadmium  in  nitric  acid,  neutralize 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  115 

by  sodium  carbonate,  precipitate  by  potassium  cyanide, 
avoiding  any  excess,  wash  the  precipitate,  and  then 
dissolve  it  in  potassium  cyanide  adding  a  slight  excess 
of  the  latter. 

Make  up  500  c.c.  of  the  bath,  and  compare  the  deposit 
on  brass  with  a  silver  deposit,  as  to  hardness,  smoothness 
and  tarnishing  in  illuminating  gas. 

• 
EXPERIMENT  99 

DEPOSITION  OF  A  COBALT-NICKEL  ALLOY 

The  alloy  containing  twenty- five  percent  cobalt  is  hard- 
est of  the  entire  cobalt-nickel  series  of  alloys  and  should  be 
especially  valuable  for  facing  electrotypes  from  which  a 
great  number  of  impressions  are  desired,  or  as  a  substitute 
for  copper  in  the  production  of  the  entire  electrotype, 
provided  thick  deposits  of  the  alloy  can  be  produced. 

Give  a  preliminary  report  on  a  suitable  electrolyte. 
Make  up  500  c.c.  of  the  approved  electrolyte,  deposit  on 
copper  in  series  with  a  nickel  bath,  and  compare  the 
deposits.  Can  thick  deposits  of  the  alloy  be  formed 
without  peeling? 

EXPERIMENT  100 

DEPOSITION  OF  SILVER  FROM  NITRATE  AND  FROM  CYANIDE 
SOLUTIONS 

Make  up  200  c.c.  of  a  solution  of  silver  nitrate  con- 
taining the  same  amount  of  metal  as  an  equal  volume  of 
the  laboratory  silver-plating  bath.  Plate  on  copper 
with  the  two  baths  in  series. 

EXPERIMENT  101 

SILVER  PLATING 

Plate  with  silver  for  fifteen  or  twenty  minutes  some 
small  article  of  your  own,  or  a  small  sheet  of  brass. 


116  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

Compare  the  appearance  of  the  deposit  after  one  minute 
with  that  when  finished.  Rinse  with  distilled  water, 
saving  the  wash  water  to  be  added  to  the  plating  bath. 
Burnish  or  polish  the  deposit. 

EXPERIMENT  102 
THE  DEPOSITION  OF  IRON 

Set  up  a  cell  containing  500  to  600  c.c.  of  the  iron 
bath,  and  with  anodes  of  flat  bar  iron  or  mild  steel, 
plate  on  brass  or  copper  for  a  half  hour,  examining  the 
deposit  occasionally.  Test  the  final  deposit  for  flexi- 
bility. Finally  put  in  a  cathode  of  clean  sheet  iron  and 
run  the  cell  for  a  week  or  two.  The  electrodes  should 
clear  the  bottom  of  the  jar  by  an  inch  to  allow  for  the 
settling  of  slime,  and  the  electrolyte  should  be  stirred 
once  or  twice  daily.  Why?  If  the  deposit  is  thicker 
upon  one  part  of  the  cathode  than  on  another,  change 
the  position  of  the  electrodes  to  secure  a  more  uniform 
distribution  of  the  current.  The  deposit  will  be  useful 
as  an  anode  for  the  production  of  iron  of  exceptional 
purity. 

EXPERIMENT  103 
DEPOSITION  OF  CADMIUM  FROM  A  CYANIDE  ELECTROLYTE 

Dissolve  25  g.  of  cadmium  chloride,  CdCl2-2H20, 
in  200  c.c.  of  water,  add  10  g.  dry  potassium  or  sodium 
carbonate,  stir  in  potassium  cyanide  solution  until  the 
precipitate  is  completely  dissolved,  and  make  up  to  500 
c.c.  with  distilled  water.  For  the  cyanide  solution  above, 
dissolve  40  g.  in  200  c.c.  water,  and  note  the  amount 
required  to  clear  up  the  precipitate.  Keep  the  remainder 
to  add  later  if  needed  to  improve  the  anode  corrosion. 

Plate  for  fifteen  minutes  on  polished  metal  at  what 
you  consider  a  suitable  current  density.  Is  the  deposit 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  117 

smooth  or  rough  in  comparison  with  similar  deposits  of 
copper,  nickel  or  silver?  Determine  the  critical  current 
density,  i.e.  that  at  which  the  deposit  becomes  rough  or 
dark.  Is  anode  corrosion  good  or  poor?  What  is  the 
current  density  at  the  anode?  If  necessary,  use  a  larger 
anode  surface  or  add  more  cyanide  until  satisfactory 
corrosion  is  secured.  Record  ratio  of  anode  to  cathode 
area,  and  total  amount  of  potassium  cyanide  required 
for  good  anode  corrosion. 

Determine  the  current  efficiency — see  experiment  90, 
What  is  the  valence  of  cadmium  in  this  solution  ?  On  what 
common  metals  does  cadmium  deposit  by  immersion? 
Do  you  think  cadmium  would  be  a  good  or  a  poor 
protective  coating  for  iron?  Compute  the  concentra- 
tion of  the  bath  in  grams  of  metal,  and  in  gram- 
equivalents  per  liter. 

EXPERIMENT  104 

THE  "  THRO  WING"  OF  COPPER 

It  is  always  a  matter  of  more  or  less  difficulty  to  secure 
a  good  deposit  of  metal  in  the  depressions  of  an  article 
to  be  plated,  and  electrolytes  differ  widely  in  their  ability 
to  deposit  metal  in  such  depressions,  or  in  their  power 
of  "  thro  wing  the  metal"  as  platers  say. 

Test  the  "  thro  wing"  of  copper  deposits  from  the  acid 
and  from  the  alkaline  bath. 

Fold  a  piece  of  paper  over  the  outside  of  a  glass  funnel 
2  1/2  or  3  inches  in  diameter,  and  cut  it  to  fit.  With  this 
as  a  model,  cut  two  sheets  of  copper,  measure  their  thick- 
ness at  several  points  with  a  micrometer,  and  fold  them 
to  the  form  of  a  cone.  If  small  projections  are  left  in 
cutting,  these  may  be  folded  over  to  hold  the  cone  to- 
gether. Deposition  on  the  outside  should  be  prevented 
by  a  coat  of  paraffine,  or  a  shield  of  paraffined  paper. 


118  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

The  anode  should  be  placed  in  front  of  the  opening  of 
the  cone  and  6  inches  or  more  from  it.  Plate  for  an 
hour  or  more  at  suitable  current  densities,  and  chart  the 
distribution  of  current  over  the  inside  of  each  funnel. 

Metallochromes 

Brilliant,  multi-colored  deposits  may  be  obtained  at 
the  anode  from  solutions  of  certain  lead  or  manganese 
salts. 

EXPERIMENT  105 
METALLOCHROMES  FROM  SODIUM  PLUMBATE 

Dissolve  5  g.  lead  nitrate  in  250  c.c.  of  water,  and  mix 
with  a  solution  of  50  g.  of  sodium  hydroxide  in  an  equal 
amount  of  water;  or  formula  No.  39  may  be  used.  The 
colors  are  most  brilliant  on  a  mirror-like  surface. 

Polish  a  sheet  of  copper  or  brass,  flash  it  for  thirty  to 
forty  seconds  in  the  nickel  bath,  and  suspend  as  anode 
in  the  above  electrolyte.  Determine  the  current  density 
and  time  required  for  the  most  satisfactory  results.  By 
using  the  point  of  a  wire  as  cathode,  at  about  1  cm. 
from  the  anode,  colored  rings  may  be  produced.  Try 
stars  and  circles  of  wire  as  cathode. 

Oxidizing  and  Coloring  Metals 

EXPERIMENT  106 
BLACK  NICKEL.     FORMULA  No.   12,  PAGE  57 

This  is  widely  used  for  the  production  of  a  durable 
black  finish  on  polished  brass  and  copper.  Polish  sheets 
of  aluminum,  brass,  copper,  iron  and  zinc,  clean  and 
plate  until  a  good  color  is  obtained.  If  deposition  occurs 
by  immersion,  use  a  high  current  for  a  few  seconds  to 
" strike"  the  work.  Will  the  deposit  stand  gentle 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  119 

polishing?     Bending?     The  durability  of  the    finish  is 
increased  by  lacquering. 

EXPERIMENT  107 

ANTIQUE  GREEN  ON  BRASS.     FORMULA  No.  30,  PAGE  80 
This  finish  is  produced  without  the  use  of  the  current. 

EXPERIMENT  108 

BLACK  ON  COPPER.     FORMULA  No.  35 
Test  the  durability  of  this  finish  as  in  experiment  106. 

EXPERIMENT  109 

BLACK  ON  COPPER.     FORMULA  No.  36 
Compare  this  finish  with  those  of  experiments  106  and  108. 

EXPERIMENT  110 
OXIDIZED  SILVER.     FORMULA  No.  38 

Silver  plate  a  sheet  of  polished  copper,  oxidize  it  in 
this  solution,  and  re-polish. 

.Plate  a  sheet  of  polished  copper  for  ten  minutes  in  the 
acid  copper  bath  to  produce  a  matte  surface,  plate  with 
silver  and  oxidize. 

EXPERIMENT  111 

GALVANOPLASTY 

This  is  the  reproduction  of  objects  in  metal  by  electro- 
deposition.  The  steps  in  the  process  are : 

1.  The  preparation  of  a  cast  from  the  object,  usually 
in  wax,  gutta-percha  or  plaster  of  Paris.     The  latter 
must  be  waterproofed  by  paraffine  or  shellac. 

2.  Rendering  the  surface  a  conductor  by  polishing  with 


120  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

finely  pulverized  graphite,  or  by  brushing  with  a  solution 
of  silver  nitrate,  followed  by  a  solution  of  potassium  sul- 
phide, or  fuming  in  hydrogen  sulphide  to  form  a  film 
of  silver  sulphide. 

3.  Deposition  of  a  thin  sheet  of  copper  upon  the  pre- 
pared surface,  requiring  ten  to  twenty  hours  in  the  acid 
copper  bath.  The  metal  shell  is  then  removed,  and 
may  be  filled  with  melted  lead  or  solder,  and  a  light 
deposit  of  any  desired  metal  may  be  given  to  the  whole. 

For  details  of  the  process,  consult  Langbein  or 
Pfanhauser. 

By  this  method  reproduce  a  medal  or  some  similar 
object.  When  finished,  it  may  be  oxidized,  and  the 
color  buffed  off  the  high  portions  so  that  they  stand  out 
against  the  darker  background. 

EXPERIMENT  112 
THE  ELECTROPLATING  OF  WOOD,  LEAVES,  ETC. 

Give  the  object  two  coats  of  very  thin  shellac,  treat 
with  silver  nitrate  and  hydrogen  sulphide  as  above  and 
plate  in  the  copper  bath.  It  may  be  necessary  to  repeat 
the  process  for  the  formation  of  silver  sulphide  several 
times  before  attempting  to  plate  with  copper.  Only 
simple  objects  like  oak  or  holly  leaves  should  be  at- 
tempted at  first. 

Electrolytic  Preparations 

References : 

Allmand — Applied  Electrochemistry,  pages  116-135,  144-147, 

386-406. 

Elbs— Electrolytic  Preparations,  pages  6-15,  89-90. 
Lehfeldt — Electrochemistry. 
Perkin — Practical  Methods  of  Electrochemistry,    pages    82, 

195-199,  250-251. 
Thompson — Applied  Electrochemistry,  pages  67-79. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  121 


Oxidation  and  Reduction 

Allmand,  pages  129-132. 

Thompson,  pages  73-74. 

Electrolysis  lends  itself  especially  well  to  oxidation 
and  reduction  processes,  since  it  is  possible  to  vary  not 
only  the  speed,  but  also  the  intensity  of  the  action  with 
great  nicety.  This  often  permits  the  reduction  of  one 
substance  in  the  presence  of  a  second  reducible  compound. 
Factors  affecting  the  intensity  of  the  reducing  action  *are 
the  material  of  the  electrode,  the  nature  of  its  surface, 
and  the  current  density.  In  comparing  the  effects  of 
different  cathodes,  an  attempt  is  frequently  made  to 
resolve  the  reducing  action  of  electrodes  into  the  catalytic 

TABLE  2. — OVERVOLTAGE  OF  HYDROGEN 


Cathode 

By 
Cas- 
par! 7 
NH2- 

SO4 

0.78 

By  Foerster  anc 
N-HzSO 

il  Si! 

ft    °    1 

Piguet8 

4 

y 

S2f 

<N      3 

By  Tafel9  0.1 
amp.  /cm.2 

Discharg 
N-l 

From 
Caspar! 

e  potentials 

12S04 

From 
Foerster 

Mercury 

0.43 

1.25 

1.32 

1.30     +.5476    +.1976 

Zinc 
Lead 

0.70 
0.64 

0  .  35 

1.26 

1.35 

1.30 

+  .4676 
+  .4076 

+  .1176 

Tin 

0.53 

0.43 

1.08 

1.16 

1.15 

+  .2976 

+  .  1676 

Cadmium 
Palladium  . 
Copper 

0.48 
0.46 
0.23 

0.48 

1.18 

1.23 

1.22 

+  .2476 
+  .2276 
-.0024 

+  .  1976 

0.10 

0.67     0.79 

0.79 

-.1324 

Nickel 
Silver 
Platinum 
Gold 

0.21 
0.15 
0.09 
0.02 

0.10 

0.07 
0.055 

0.64 

0.74  J0.74 
0.93? 

-.0224 
-.0824 
-  .  1424 
-.2124 

-.1324 

-  .  1624 

-  .  1874 

0.86 

0.96    0.95 

Platinized- 

0.0 

O.OOSi  0.05 

C.07  10.07 

-.2324!    -.2274 

platinum. 

7Z.  phys.  Chem.,  1899,  30,  89. 
8Z.  f.  Elektrochem.,  1904,  10,  715. 
9Z.  phys.  Chem.  1904,  50,  712. 


122  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

action  of  the  electrode  material,  and  what  is  called  the 
overvoltage  of  hydrogen.  Current  density  is  supposed 
to  affect  only  the  latter. 

The  variation  in  the  potential  required  by  electrodes 
of  different  metals  for  visible  evolution  of  hydrogen  is 
usually  expressed  as  the  "overvoltage  of  hydrogen" 
on  that  particular  metal,  the  least  potential  of  platinized 
platinum  for  hydrogen  evolution  being  taken  as  zero. 
Table  2  gives  values  of  the  overvoltage  of  hydrogen 
obtained  by  different  observers.  The  effect  of  current 

TABLE  3. — ANODE  POTENTIALS  AND  OVERVOLTAGE  OF  OXYGEN 


i~&s 

CO 

C-3J2        '«  £  i 

DA  =  0.033  amp./cm«. 

1 

Anode 

fills 

c  o  vg  "3  j^ 

I||K 

Overvoltagc 
Allmand,  page 

Discharge  pot( 
tial  vs.  calom 
electrode  calc. 
author 

By  Foerster 
Least  potential 
evolution  hyd. 
hyd.  electrod 
2NKOH 

£0     ! 

£o              M- 

o3  LO            PH    . 
HTH^      !      OO 

HH       -          1          K>  O 

SI      §§ 

*~?   0               (N 
(NIM 

2N-H2SO4, 
99°C. 

Nickel, 

1.28 

0.05    -0.9524 

sponge. 

Nickel, 

1.35 

0.12 

-1.0224 

1.35 

2.00     1.77 

smooth. 

Cobalt  .    .  . 

1.36 

0.13 

-1  0324 

Iron 

1  47 

0  24 

-1  1424 

1  47 

2  02     1  89 

Platinized 

1.47  JO.  24 

-1.1424 

1.47 

2  .  30    

Pt, 

Copper. 

1  48   0  25 

-1  1524 

Lead  

1.53  10.30 

-1.2024    

Silver  

1.63    0.40 

-1.3024  i  

Cadmium  . 

1.65 

0.42 

-  1  .  3224 

Palladium  . 

1.65    0.42 

-1.3224 

1.65 

2.45    

Platinum  .  . 

1.67    0.44    -1.3424  j     1.67 

2.92     2.50  i  2.17 

Gold  

1.75    0.52    -1.4224    

density     is     also     shown.     The     discharge     potentials, 
referred   to   the    calomel   electrode  (value   —0.56  volt) 


10  Coehn  and  Osaka — Z.  Anorg.  Chem.,  1903,  34, 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY   123 

have  been  calculated  by  the  author  from  the  values  for 
overvoltage  by  the  use  of  Wilsmore's  value  0.3276  volts 
for  the  difference  between  the  calomel  electrode  and  the 
hydrogen  electrode  in  normal  sulphuric  acid.11 

Oxygen  shows  similar  overvoltage  effects  upon  anodes 
of  different  materials,  as  shown  in  table  3. 

The  increase  of  overvoltage  with  time  and  its  diminu- 
tion with  rise  of  temperature  varies  for  different  metals. 


EXPERIMENT  113 

THE  EFFECT  OF  CATHODES  OF  DIFFERENT  METALS  ON 
THE  REDUCTION  OF  POTASSIUM  NITRATE 

Text  the  reducing  action  on  normal  potassium  nitrate 
solution  of  cathodes  of  platinized  platinum,  smooth  plat- 
inum, and  polished  zinc.  These  experiments  may  most 
conveniently  be  carried  out  by  connecting  a  standard  oxy- 
hydrogen  coulombmeter  in  series  with  the  same  appara- 
tus (see  experiment  25)  containing  the  special  cathode. 
If  a  coulombmeter  is  not  available,  a  cell  may  be  made 
from  two  wide  mouth  bottles  fitted  with  two-hole 
rubber  stoppers.  One  hole  of  each  carries  a  short  glass 
tube  connecting  the  bottles,  and  through  the  other  hole 
passes  a  glass  tube  connected  to  a  gas  burette  or  other 
device  for  collecting  gases.  The  electrodes  are  carried 
by  stout  wires  passing  through  the  stoppers. 

By  means  of  the  gas  coulombmeter,  or  an  accurate 
ammeter  placed  in  series  with  the  experimental  cell, 
determine  the  hydrogen  equivalent  of  the  current  used. 
From  this  and  the  volume  of  gas  collected,  determine 
what  percent  of  the  current  was  spent  in  reduction  with 
each  cathode. 

"Lehfeldt,  page  240. 


124  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

EXPERIMENT  114 

THE    OXIDATION    OF    LITHARGE    TO    LEAD    PEROXIDE. 
PERKIN,  PAGE  215 

The  electrolyte,  consisting  of  a  twenty  percent  solu- 
tion of  sodium  chloride,  should  be  placed  in  a  rectangular 
battery  jar,  with  an  anode  of  platinum  or  graphite  and  a 
cathode  of  lead  or  graphite.  The  current  density  at  the 
anode  may  be  1  to  1.5  amperes.  The  cathode  should  be 
wrapped  in  parchment  or  cloth  (why?)  and  the  electrolyte 
must  be  stirred  vigorously  during  electrolysis  to  keep 
the  litharge  in  suspension.  Use  25  g.  of  finely  pulverized 
litharge.  This  becomes  darker  in  color*  as  electrolysis 
proceeds,  -and  when  it  is  all  converted  to  the  peroxide,  the 
color  is  a  dark  brown. 

After  passing  the  current  for  the  number  of  hours  re- 
quired by  theory,  collect  the  product  on  a  filter,  wash  it 
free  from  chloride,  and  warm  it  with  dilute  nitric  acid  to 
remove  any  unaltered  litharge,  wash,  dry  and  weigh. 
Ascertain  the  total  yield,  grams  per  ampere-hour  and  the 
current  efficiency.  No  chlorine  is  liberated  during  elec- 
trolysis, since  sodium  hypochlorite  is  formed,  and  it  is 
this  mainly  which  oxidizes  the  litharge. 

EXPERIMENT  115 
LEAD  PEROXIDE  FROM  LEAD  NITRATE 

Dissolve  38  g.  lead  nitrate  in  500  c.c.  of  water,  and 
12  g.  of  caustic  soda  in  400  c.c.  of  water.  Mix  the  solu- 
tions, add  100  g.  of  salt  and  2  g.  of  potassium  chromate, 
start  the  stirrer,  electrolyze  and  purify  as  before.  The 
purpose  of  the  chromate  is  to  form  a  diaphragm  of 
chromium  hydrate  over  the  cathode,  and  so  lessen  reduc- 
tion of  the  hypochlorite  by  nascent  hydrogen.  The 
cathode  need  not  be  covered  with  parchment  in  this  ex- 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY   125 

periment.  Compare  the  yield  and  current  efficiency  with 
the  results  in  the  previous  experiment.  To  what  do 
you  ascribe  the  difference?  To  insure  the  same  current 
in  both  experiments,  they  may  be  run  in  series. 

EXPERIMENT  116 

LEAD    PEROXIDE   FROM   LEAD   BY  LUCKOW'S  PROCESS. 
ALLMAND,  PAGE  389 

Use  weighed  electrodes  of  sheet  lead  and  an  electrolyte 
containing  15  g.  per  liter  of  a  mixture  of  99.5  parts  so- 
dium sulphate  and  0.5  parts  sodium  chlorate,  slightly 
acidified  with  sulphuric  acid.  Determine  the  current 
efficiency  and  the  actual  yield.  What  are  the  defects  of 
the  process?  The  effect  of  different  current  densities 
may  be  investigated. 

EXPERIMENT  117 

POTASSIUM  PERSULPHATE.     ELBS,  PAGE  138; 
PERKIN,  PAGE  202 

This  preparation  requires  that  the  temperature  be  kept 
below  20°  C.  and  a  better  yield  is  obtained  below  10°  C. 
The  electrolyte  consists  of  a  saturated  solution  of  po- 
tassium bisulphate,  the  strength  of  which  is  maintained 
by  suspending  crystals  of  KHS04  in  a  perforated  vessel 
in  the  upper  part  of  the  electrolyte.  The  anode  consists 
of  a  spiral  of  platinum  wire  of  1  to  2  sq.  cm.  surface, 
sealed  into  a  glass  tube  and  placed  near  the  bottom  of  the 
electrolyte.  This  should  be  surrounded  by  a  wide  glass 
tube  open  at  both  ends,  and  the  cathode  of  platinum  or 
lead  should  surround  this  tube  near  the  surface  of  the 
electrolyte,  which  is  contained  in  a  very  tall,  narrow 
beaker,  placed  in  a  large  dish  of  ice  water.  The  current 
density  at  the  anode  should  be  between  500  and  1000 


126  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

amperes  per  sq.  dm.  and  that  at  the  cathode  as  low  as 
possible  to  minimize  heating.  Mueller  finds  that  the 
yield  of  persulphate  is  increased  by  adding  small  amounts 
of  hydrofluoric  acid;  this  requires  the  use  of  hard  rubber 
or  paraffined  glass  apparatus. 

As  soon  as  the  electrolyte  becomes  saturated  with  the 
slightly  soluble  persulphate,  the  solid  salt  begins  to  sepa- 
rate. After  a  considerable  amount  has  formed,  collect 
and  dry  the  product  with  the  aid  of  a  filter  pump,  and 
complete  the  drying  over  sulphuric  acid  in  vacuo.  Deter- 
mine its  purity,  and  the  current  efficiency. 

The  Estimation  of  Persulphuriq  Acid 

To  the  solution  of  persulphuric  acid,  add  a  known 
amount  of  a  solution  of  ammonio-ferrous  sulphate,  and 
titrate  the  excess  of  this  by  permanganate. 

H2S208  +  2FeS04  =  Fe2(S04)3  +  H2S04 

For  persulphates,  the  addition  of  ferrous  sulphate  should 
be  followed  by  one-half  volume  of  dilute  sulphuric  acid 
and  100  c.c.  of  boiling  water,  and  the  solution  should  be 
titrated  by  permanganate  at  once.  Potassium  or  sodium 
persulphate  may  be  estimated  by  determining  the  loss  of 
weight  on  ignition. 

K2S208  =  K2S04  +  S03  +  0 

The  preparation  of  ammonium  presulphate  requires  a 
diaphragm,  but  is  more  satisfactory  than  the  previous 
experiment.  A  diaphragm  may  be  used  in  the  prepara- 
tion of  potassium  persulphate  with  good  results. 

EXPERIMENT  118 

AMMONIUM  PERSULPHATE.     ELBS,  PAGE  35; 

PERKIN,  PAGE  204 

The  electrodes,  current  density  and  temperature 
are  as  in  the  previous  experiment,  except  that  the  cathode 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY   127 

may  be  a  lead  tube  carrying  cold  water.  A  porous  cup, 
boiled  for  several  hours  in  water  to  expel  the  air,  serves 
as  a  diaphragm  and  contains  the  anolyte — the  solution 
surrounding  the  anode — which  consists  of  2  percent 
sulphuric  acid  saturated  at  10°  C.  with  ammonium 
sulphate.  The  strength  of  the  anolyte  is  maintained 
as  in  experiment  117.  To  diminish  its  resistance,  the 
porous  cup  should  be  allowed  to  soak  in  the  anolyte 
over  night.  The  catholyte  consists  of  1  volume  of 
concentrated  sulphuric  acid  to  1  1/2  volumes  of  water. 
The  anolyte  may  be  regenerated  after  use  by  the  addition 
of  ammonia  saturated  with  ammonium  sulphate,  the 
catholyte  by  the  addition  of  sulphuric  acid.  Proceed  with 
the  electrolysis,  collection  of  the  product  and  titration  as 
in  experiment  117. 

EXPERIMENT  119 

POTASSIUM    CHLORATE    FROM    POTASSIUM    CHLORIDE. 

ELBS,  PAGE  26,  29-31;  PERKIN,  PAGE  210; 

HOSTELET,  PAGE  79 

The  electrolyte  consists  of  100  g.  potassium  chloride, 
1  g.  potassium  carbonate,  and  1  g.  potassium  dichromate 
dissolved  in  250  c.c.  of  warm  water.  The  anode  should 
be  platinum  gauze  or  sheet,  and  the  cathode  a  sheet  of 
platinum,  nickel,  copper  or  graphite.  The  electrodes 
should  be  placed  1  cm.  apart.  The  current  density 
at  the  anode  should  be  20  amperes  per  sq.  dm., 
but  at  the  cathode  should  be  greater.  Why?  Dur- 
ing electrolysis  the  electrolyte  should  be  maintained 
at  a  temperature  of  50°  to  60°  C.  kept  faintly  acid  by  a 
stream  of  carbon  dioxide,  and  gently  stirred.  Air  may 
be  used  for  stirring. 

6KOH  +  6C1  =  KC103  +  5KC1 
One  ampere-hour  liberates  1.322  g.  of  chlorine.    Compute 


128  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

the  amount  of  potassium  chlorate  which  should  be  formed 
per  ampere-hour.  Pass  100  to  400  ampere-hours, 
using  a  coulombmeter  or  recording  ampere-hour  meter 
in  series  with  the  cell.  Potassium  chloride  should  be 
supplied  from  a  bag  or  perforated  cup  hung  in  the  upper 
part  of  the  cell. 

After  cooling,  collect  the  potassium  chlorate,  recrystal- 
lize  it  once,  and  weigh  it.  Evaporate  the  mother  liquor 
from  the  cell  to  half  its  volume  and  cool  it.  Recrystallize 
this  product  also.  Instead  of  evaporating  the  contents 
of  the  cell,  the  amount  of  dissolved  chlorate  may  be 
determined  by  volumetric  analysis.  For  the  solubilities 
of  potassium  chloride  and  potassium  chlorate,  see  Elbs, 
page  27.  State  the  current  efficiency  based  on  the  total 
amount  of  chlorate  formed.  It  should  exceed  70  percent 
until  over  half  the  chloride  has  been  converted  to 
chlorate. 

EXPERIMENT  120 

PREPARATION  OF  OTHER  CHLORATES 

The  chlorates  of  sodium,  barium,  calcium  and  stron- 
tium may  be  prepared  in  a  manner  similar  to  that  for 
potassium  chlorate,  but  since  these  chlorates  are  very 
soluble,  it  is  difficult  to  separate  them  from  the  chlorides. 
Prepare  barium  or  sodium  chlorate.  How  does  the  cur- 
rent efficiency  compare  with  that  of  potassium  chlorate? 
How  could  you  make  the  latter  from  your  electrolyte? 
How  can  you  prepare  a  solution  of  copper  chlorate  from 
your  product? 

EXPERIMENT  121 

POTASSIUM  BROMATE  FROM  POTASSIUM  BROMIDE.     ELBS, 
PAGE  28;  PERKIN,  PAGE  212 

The  electrolyte  consists  of  125  g.  potassium  bromide, 
and  1  g.  potassium  chromate  in  500  c.c.  of  water.  What 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  129 

is  the  purpose  of  adding  the  chromate  ?  The  anode  should 
be  platinum  gauze  or  sheet  and  the  cathode  a  sheet  of 
platinum  or  nickel.  The  current  density  at  the  electrodes 
should  be  between  10  and  12  amperes  per  sq.  dm.  and 
the  temperature  40°  C. 

6KOH  +  6Br  =  KBr03  +  5KBr 

One  ampere-hour  liberates  2.982  g.  of  bromine.  Calculate 
the  amount  of  bromate  that  should  be  produced  per 
ampere-hour.  Pass  125  to  150  ampere-hours,  evaporate 
to  0.4  volume,  and  cool.  Collect  and  dry  the  bromate 
by  aid  of  the  filter  pump.  The  current  efficiency  should 
be  ninety  percent,  -and  the  material  yield  about  seventy 
percent  of  the  theoretical  amount. 
100  parts  of  water  dissolves 

3.1  parts  potassium  bromate  at  0°  C. 

7.0  parts  potassium  bromate  at  20° 
22.8  parts  potassium  bromate  at  60° 
49.8  parts  potassium  bromate  at  100° 

EXPERIMENT  122 

POTASSIUM  IODATE  FROM  POTASSIUM  IODIDE 

Use  as  electrolyte  25  g.  potassium  iodide,  1  g.  potassium 
hydrate  and  0.2  g.  potassium  chromate  per  100  c.c. 
at  25°  to  30°  C.,  with  other  conditions  as  for  potassium 
bromate. 

100  parts  of  water  dissolves 

4.7  parts  potassium  iodate  at      0°  C. 

8.1  parts  potassium  iodate  at    20° 
18.5  parts  potassium  iodate  at    60° 
32.2  parts  potassium  iodate  at  100° 

Sodium  iodate  is  prepared  similarly  and  is  only  slightly 
more  soluble.  Which,  then,  can  probably  be  prepared 
at  the  higher  current  efficiency?  Why? 


130  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

EXPERIMENT  123 

POTASSIUM   PERCHLORATE.     ELBS,   PAGE    33;    PERKIN, 
PAGE  214. 

The  electrolyte  consists  of  a  saturated  solution  of 
potassium  chlorate,  with  a  perforated  vessel  or  bag  con- 
taining crystals  of  the  same  salt  suspended  near  the  top 
of  the  cell.  The  anode  consists  of  platinum  gauze  or 
sheet,  and  the  cathode  of  a  sheet  of  platinum  or  copper. 
The  current  density  at  the  anode  should  be  less  than  at 
the  cathode,  and  may  be  from  8  to  12  amperes  per 
sq.  dm.  The  temperature  must  be  kept  below  25° 
C.,  and  the  efficiency  is  better  if  a  temperature  of  10°  C. 
is  maintained.  Occasional  but  not  continuous  stirring 
is  advantageous.  A  current  efficiency  of  about  eighty 
percent  may  be  attained. 

100  parts  of  water  dissolves 

0.7  parts  potassium  perchlorate  at  0°  C. 

6.4  parts  potassium  perchlorate  at  50° 

19.9  parts  potassium  perchlorate  at  100° 


Potassium  perchlorate  may  be  prepared  with  a 
current  efficiency  of  ninety  percent  by  electrolysis  of  a 
solution  containing  per  liter  300  to  500  g.  of  sodium 
chlorate,  and  at  the  end  of  the  experiment  adding  a 
quantity  of  a  cold  saturated  solution  of  potassium 
chloride  corresponding  to  the  ampere-hours  passed. 

NaC103  +  O  =  NaClO4 

Each  100  g.  of  perchlorate  requires  61  g.  potassium 
chloride,  or  190  c.c.  of  a  solution  of  the  chloride  saturated 
at  20°  C.  Should  the  electrolyte  become  alkaline  during 
electrolysis,  it  should  be  made  faintly  acid  by  the  addi- 
tion of  a  few  drops  of  dilute  sulphuric  acid. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  131 

EXPERIMENT  124 
BARIUM  PERCHLORATE.     ELBS,  PAGE  34 

Electrolyze  as  in  experiment  123  a  cold  solution  con- 
taining 300  g.  barium  chlorate,  BaC103-H2O  per  liter. 
The  current  efficiency  exceeds  seventy  percent  at  the  out- 
set, but  drops  to  twenty  percent  after  ninety  percent  of 
the  chlorate  has  been  changed  to  perchlorate.  What 
should  be  the  effect  of  stirring  at  the  anode  at  this 
stage  of  the  electrolysis?  Since  both  salts  are  very 
soluble,  a  separation  of  the  two  in  aqueous  solution 
is  not  practicable.  Evaporate  to  dryness  on  a  water 
bath,  and  extract  the  residue  with  hot  ninety-five  percent 
alcohol,  in  which  the  perchlorate  is  readily  soluble, 
but  the  chlorate  and  chloride  are  almost  insoluble. 
Cautiously  distil  off  the  alcohol.  Mixtures  of  com- 
bustible substances  with  chlorates  or  perchlorates  may 
be  exploded  by  heat  or  percussion.  Strong  and  sud- 
den heating  of  perchlorates  should  be  avoided  as  several 
of  them  are  explosive. 

Barium  perchlorate  is  especially  useful  for  the  prepara- 
tion of  perchlorates  of  other  metals  for  use  in  electro- 
plating, for  which  purpose  the  perchlorates  are  proving 
particularly  advantageous.  Outline  a  method  for  pre- 
paring a  solution  of  copper  perchlorate  from  barium 
perchlorate. 

EXPERIMENT  125 

OXIDATION  OF  CHROMIUM  SULPHATE  TO  CHROMIC 
ACID.     ELBS,  PAGE  17 

Use  as  anode  a  sheet  of  lead  previously  per-oxidized 
by  use  as  anode  in  dilute  sulphuric  acid.  The  cathode  of 
sheet  lead  should  be  placed  in  a  porous  cup  which  is 
surrounded  by  the  anode.  The  anolyte  consists  of 
150  c.c.  of  concentrated  sulphuric  acid,  200  g.  of  chrome 


132  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

alum,  KCr (804)2*  12  H2O,  per  liter.  This  is  prepared 
by  pouring  the  acid  slowly  into  about  750  c.c.  of  water, 
and  stirring  in  the  powdered  alum.  The  catholyte  may 
consist  of  one  volume  of  sulphuric  acid  to  four  volumes 
of  water.  The  temperature  should  be  from  50°  to  60°  C. 
and  the  current  density  at  the  anode  2  to  3  amperes  per 
sq.  dm.  How  could  you  raise  the  current  density 
at  the  anode  without  lessening  the  current  effi- 
ciency? What  should  be  the  current  density  at  the 
cathode? 

Cr2(SQ4)3  +  3(S04)+  6H2O  =  2Cr03  +  6H2SO4 
Stir  every  half  hour  and  draw  out  a  sample  for  titra- 
tion,  which  can  be  performed  by  adding*  excess  of  care- 
fully weighed  ferrous  ammonium  sulphate,  and  titrat- 
ing with  permanganate.  The  theoretical  yield  is  1.25  g. 
CrOs  per  ampere-hour.  Maintain  the  current  constant, 
and  plot  curves  of  current  efficiency  and  of  grams  CrO3 
per  liter  vs.  time.  The  current  efficiency  should  exceed 
ninety  percent  until  most  of  the  chromium  salt  has  been 
oxidized. 

EXPERIMENT  126 

CUPROUS  AND  CUPRIC  HYDROXIDES   OR  OXIDES  FROM 
COPPER.     ELBS,  PAGE  40;  PERKIN,  PAGE  220 

Arrange  two  cells  in  series,  one  containing  a  thirteen 
percent  solution  of  sodium  chloride,  the  other  a  sixteen 
percent  solution  of  crystallized  sodium  sulphate.  The 
copper  anodes  and  iron  cathodes  should  be  4  cm.  apart, 
and  the  same  distance  from  the  bottom  of  the  cells. 
The  anodes  should  be  enclosed  in  cloth  bags  or  wrapped 
in  parchment.  Why?  The  electrolyte  should  be  stirred. 
At  room  temperature  the  hydroxides  are  formed,  but 
at  100°  C.,  the  products  are  the  oxides.  Wash,  dry 
and  weigh  the  oxides,  determine  the  anode  corrosion, 
the  current  efficiency  at  the  anode,  and  the  current 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  133 

efficiency  of  the  process.  What  would  be  the  cost  of 
electric  energy  at  10  cents  per  kilowatt  hour  for  a 
pound  of  each  product  according  to  the  conditions  of  this 
experiment  ?  What  data  do  you  need  besides  the  ampere- 
hours  per  pound  of  product  in  order  to  answer  the  last 
question?  See  that  the  necessary  data  are  obtained 
during  the  experiment. 

EXPERIMENT  127 
POTASSIUM  PERMANGANATE.     PERKIN,  PAGE  218 

Manganese  or  ferro-manganese  serves  as  anode  in  a 
forty  percent  solution  of  potassium  carbonate,  specific 
gravity  1.42.  The  cathode  of  sheet  iron  must  be  placed 
in  a  porous  cup.  Iron  is  precipitated  as  hydroxide. 
The  current  efficiency  may  be  followed  by  titrating  with 
oxalic  acid. 

2KMn04  +  5H2C2O4+   3H2SO4  =  10CO2  +  K2S04 
+  2MnS04  +  8H20. 

Find  the  anode  corrosion  and  the  amount  of  permanga- 
nate produced.  Plot  the  curve  of  grams  of  permanganate 
against  ampere-hours  or  time.  If  a  plate  of  copper 
oxide  or  the  positive  plate  of  a  storage  cell  be  substituted 
for  the  iron  cathode,  the  porous  cup  may  be  omitted. 
Why  is  this?  The  yield  is  poor. 

EXPERIMENT  128 
POTASSIUM  CHROMATE.     PERKIN,  PAGE  219 

By  the  method  of  the  previous  experiment,  potassium 
chromate  may  be  made  by  using  an  anode  of  ferro- 
chromium  and  potassium  hydrate  as  electrolyte.  The 
progress  of  electrolysis  may  be  followed  by  occasional 
titration  of  a  sample  with  permanganate  after  the  addi- 
tion of  an  excess  of  ferrous  ammonium  sulphate  and 
sulphuric  acid. 


134  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 


2K2Cr04  +  GFeSOi  +  8H2S04  =  3Fe2(S04)3  +  Cr2 
(S04)3  +  2K2S04  +  8H20. 

The  yield  is  good.     Find  the  current  efficiency  at  the 
anode  as  well  as  of  the  production  of  chromate. 

EXPERIMENT  129 

THE  REFINING  OF  MERCURY.     HOSTELET, 
PAGE  88 

As  electrolyte  prepare  a  solution  of  mercurous  nitrate 
by  the  action  of  nitric  acid  on  excess  of  mercury.  Place 
the  mercury  to  be  purified  as  anode  in  a  small  crystalliz- 
ing dish  set  in  a  larger  one.  As  cathode  use  a  platinum 
wire  or  purified  mercury  in  the  outer  disk.  The  current 
density  at  the  anode  may  be  1  to  1.5  amperes  per  sq. 
dm.  To  prevent  the  solution  about  the  anode  from 
becoming  saturated,  it  should  be  stirred  occasionally. 
Determine  the  current  efficiency  at  anode  and  cathode. 
Assuming  that  the  impure  mercury  from  use  in  the 
laboratory  had  dissolved  traces  of  cadmium,  copper, 
gold,  silver,  and  zinc,  which  of  these  impurities  would 
still  be  found  in  the  anode,  and  which  in  the  electrolyte 
at  the  end  of  the  experiment?  Why? 

EXPERIMENT  130 

SCHEELE'S   GREEN,   CuHAsO3.     POISON! 

As  electrolyte  use  10  g.  of  crystallized  sodium  sulphate 
per  liter,  with  electrodes  of  sheet  copper  placed  far  apart, 
and  suspend  near  the  cathode  a  cloth  bag  containing 
arsenious  oxide.  Use  a  current  density  at  the  anode  of 
2.5  amperes  per  sq.  dm.,  and  stir  vigorously.  The 
products  of  electrolysis  are  copper  sulphate  and  caustic 
soda.  The  latter  dissolves  arsenious  oxide,  forming 
sodium  arsenite,  NaaAsOs,  which  reacts  with  the  copper 
sulphate  to  form  Scheele's  green.  Determine  the  loss 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  135 

of  weight  of  anode,  the  efficiency  of  anode  corrosion,  and 
the  amount  of  Scheele's  green  produced  per  ampere- 
hour.  Discuss  the  convenience  and  cost  of  the  electro- 
lytic vs.  the  chemical  manufacture  of  this  compound. 

EXPERIMENT  131 

SODIUM  HYPOCHLORITE   FROM  SALT.     ELBS,  PAGE    21; 
PERKIN,  PAGE  207 

When  a  strong  solution  of  sodium  chloride  is  electro- 
lyzed  cold  without  a  diaphragm,  the  main  product  is 
sodium  hypochlorite. 

2NaOH  +  2C1  =  NaOCl  +  NaCl  +  H20. 

As  soon  as  hypochlorite  is  present  in  considerable  amount 
it  reaches  the  cathode  and  is  reduced. 

NaOCl  +  H  =  NaCl  +  H20. 

Since  this  reduction  is  caused  only  by  nascent  hydrogen, 
a  high  current  density  at  the  cathode  is  desirable. 

With  low  current  density  at  the  anode,  there  is  prac- 
tically no  evolution  of  gas  at  first,  but  as  the  proportion 
of  hypochlorite  increases,  oxygen  is  evolved  at  the  anode, 
since  hypochlorite  is  more  readily  decomposed  than 
chloride.  With  high  current  density  at  the  anode,  the 
solution  about  the  anode  is  kept  impoverished  in  CIO 
ions,  and  this  undesirable  reaction  is  at  a  minimum. 

As  electrolyte  use  500  to  1000  c.c.  of  a  clear,  saturated 
solution  of  salt,  an  anode  of  platinum,  nickel  or  graphite. 
The  current  density  at  the  anode  should  be  12  to  16 
amperes,  and  that  at  the  cathode  20  to  30  amperes  per 
sq.  dm.  The  temperature  should  be  between  15°  and 
20°  C.  Above  25°  C.  there  is  formation  of  chlorate. 

Circulation  may  be  induced  by  placing  the  anode  on 
one  side  of  the  cell  near  the  top  and  the  cathode  on  the 
other  side  near  the  bottom.     Maintain  a  constant  known 
10 


136  A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY 

current,  and  every  half  hour  stir,  remove  10  c.c.  of  the 
electrolyte  and  titrate  by  adding  excess  of  potassium 
iodide,  acidifying  strongly  with  acetic  acid,  and  estimat- 
ing the*  free  iodine  by  tenth  normal  sodium  thiosul- 
phate.  Hydrochloric  acid  must  not  be  used  for  acidi- 
fying, as  is  often  recommended,  since  this  slowly 
decomposes  any  chlorate  present. 

Plot  curves  of  current  efficiency  against  time,  and 
against  the  percent  of  hypochlorite  formed.  For  the 
method  of  determining  the  percent  of  the  total  current 
which  is  wasted  by  reduction  and  in  the  useless  evolution 
of  oxygen  at  the  anode,  consult  either  of  the  references 
cited  above. 

The  addition  of  0.5  percent  of  sodium  or  potassium 
chromate  causes  the  formation  of  a  film  of  chromium 
hydrate  over  the  cathode,  which  greatly  lessens  the 
undesirable  reduction  of  hypochlorite.  A  second  experi- 
ment may  be  tried  with  this  addition. 

EXPERIMENT  132 

THE  OXIDATION  OF  POTASSIUM  FERROCYANIDE  TO  THE 
FERRICYANIDE.     V.  HAYEK12 

A  diaphragm  is  required  and  the  anode  compartment 
should  have  twice  the  volume  of  the  cathode  compart- 
ment. The  anolyte  should  be  well  stirred  continuously, 
and  should  be  kept  faintly  alkaline  and  at  a  temperature 
of  25°  C.  Electrodes  of  nickel  gauze  or  sheet  may  be 
used,  with  an  anode  current  density  of  0.3  to  1.5  amperes 
per  sq.  dm.,  depending  on  circulation  and  the  concen- 
tration of  the  solution,  which  may  be  from  ten  percent 
to  saturation. 

After  complete  conversion  to  the  ferricyanide,  as  shown 
by  its  no  longer  producing  a  blue  precipitate  with  a 

12  Z.  anorg.  Chem.,  1904,  39,  240. 


A  LABORATORY  COURSE  IN  ELECTROCHEMISTRY  137 

drop  of  a  solution  of  a  ferric  salt,  crystallize  and  weigh 
the  product.  Find  the  material  yield  and  the  current 
efficiency. 

EXPERIMENT  133 

POTASSIUM  FROM  POTASSIUM  HYDRATE.     HOSTELET, 

PAGE  70 

Melt  down  to  quiet  fusion  in  an  iron  crucible  400 
to  500  g.  of  potassium  hydrate.  In  about  two  hours  it 
should  be  ready  to  electrolyze.  Thrust  the  cathode  of 
iron  wire  3  mm.  in  diameter  through  a  hole  in  the  bottom 
of  a  magnesia  crucible  and  invert  this  in  the  electrolyte. 
As  anode  use  a  sheet  of  iron.  Explosions  occur  at  first, 
due  to  the  burning  of  the  potassium  in  the  air  contained 
in  the  crucible.  Fifteen  amperes  have  given  12  to  13  g. 
of  potassium  per  hour,  a  yield  of  about  fifty-eight  percent. 
On  stopping  electrolysis,  allow  the  cell  to  cool,  and  remove 
the  crucible  containing  the  potassium  only  just  before 
it  would  freeze  into  the  electrolyte,  and  plunge  it  into 
kerosene  to  cool. 


TABLE  4.- 


APPENDIX 

-ATOMIC  WEIGHTS  AND  ELECTROCHEMICAL 
EQUIVALENTS 


:  Atomic 
;    weight 

Valencel 

Grams 
per  amp.- 
hour 

1  Atomic 
i  weight 

Grams 
per  amp.  - 
hour 

Aluminum}  27.1 

3   0.3368 

Magnesium 

24.3   2 

0.4531 

Antimony 

120.2 

3 

1.4966 

Manganese 

54.9   2 

1.0255 

Arsenic 

74.9 

3 

0.9324 

Mercury 

200.6 

1 

7.4803 

Barium 

137.4 

2 

2.5619 

Molybdenum 

96.0 

2 

1.7900 

Bismuth     208.0 

3 

2.5854 

Nickel 

«.58.7 

2 

1.0945 

Bromine 

79.9 

1 

2.9814 

Nitrogen 

14.0 

3 

0.1745 

Cadmium   112.4 

2 

2.0955 

Oxygen 

16.0 

2 

0.2983 

Calcium 

40.1 

2 

0.7477 

Palladium         107  .  0 

2 

1.9951 

Carbon 

12.0 

4 

0.1118 

Platinum           195.2 

4 

1.8206 

Chlorine 

35.5 

1 

1  .  3220 

Potassium 

39.1 

1 

1.4584 

Chromium 

52.0 

3 

0.6476 

Silicon 

28.3 

4 

0.2638 

Cobalt 

59.0 

2  ll.lOOO 

Silver              !  107.9 

1 

4.0248 

Copper 

63.6 

2 

1  .  1858 

Sodium 

23.0 

1 

0.8596 

Fluorine 

19.0 

1 

0.7085 

Strontium 

87.6 

2 

1.6333 

Gold 

197.2 

3 

2.4513 

Sulphur 

32.1 

2 

0.5980 

Hydrogen 

1.008 

1 

0.03759 

Tin 

119.0 

2 

2.2188 

Iodine 

126.9 

1 

4  .  7303 

Titanium 

48.1 

4 

0.4490 

Iron 

55.8 

2 

1.0404 

Tungsten 

184.0 

2 

3.4308 

Lead          j  207.1 

2 

3.8613 

Zinc                  !    65.4 

2 

1.2194 

Lithium     f      6.9 

1 

0.2622 

26.  817    ampere-hours    deposit    one    gram-equivalent    of    any 
substance. 


138 


APPENDIX 


139 


TABLE  5. — RESISTIVITY  OF  METALS.     MICROHMS  PER  CM3. 


At  0°  C. 

Resis- 
tivity 

Temp, 
coeff. 
for  1° 
C. 

Resis- 
tivity 

Temp, 
coeff. 
for  1° 
C. 

Aluminum 
Antimony 
Arsenic 
Bismuth 

Annealed 

2.8 
36.0 
33.3 
130.0 

0.0046 
0  .  0035 

Nickel 
Ni  chrome 
Platinum 
Silver 

Annealed 
Annealed 

12.4 
95.5 
9.0 
1.50 

0.00538 
0.00043 
0.00341 
0.00377 

Copper 
Gold 
German 
silver. 
Iron 
Lead 
Mercury 

Annealed 
Annealed 

Annealed 

1.58 
2.0 
20.8 

9.5 
19.0 
94.2 

0.0039 
0.0037 
0.0004 

0.0058 
0.0038  ! 
0.00072: 

Tin 
Zinc 

10.0 
5.6 

0.00428 
0.00365 

TABLE  6. — RESISTIVITY  OF  ELECTROLYTES. 
HOLBORN 


KOHLRAUSCH    & 


H2SO*  at  18°  C. 

Grams 
.  ,  .             Specific 
acid  m 

lOOg.ofsol.1     gravity 

Resistivity 

Temp, 
coefficient 
for  1°  C. 

Gram 
equivalents 
per  liter 

1         |  

21.93 

0.00112 

0.204 

2.5 

1.0161 

9.24 

0.00115 

0.519 

5 

1.0331 

4.82 

0.00121 

1.065 

10 

1.0673 

2.57 

0.00128 

2.182 

15 

1.1036 

1.85 

0.00136 

3.384 

20 

1.1414 

1.54 

0.00145 

4.667 

30 

1.2207 

1.36 

0.00162 

7.487 

40 

1.3056 

1.48 

0.00178 

10.68 

50 

1.3984 

1.86 

0.00193 

14.30 

60 

1.5019 

2.70 

0.00213 

18.42 

70 

1.6146 

4.67 

0.00256 

23.11 

80 

1  .  7320 

9.13 

0.00349 

28.33 

85 

.7827 

10.30 

0.00365 

30.98 

90 

.8167 

9.38 

0.00320 

33.43 

95 

.8368 

9.84 

0.00279 

35.68 

97 

.8390 

12.50 

0.00286 

36.47 

99.4 

.8354 

118.00 

0.00400 

37.22 

140 


APPENDIX 


TABLE  6. — RESISTIVITY  OF  ELECTROLYTES. — Continued.     KOHL- 

RAUSCH    &    HOLBORN 


HC1  at  10° 

G.  HCl 

per  100  g. 
of  sol. 

Specific 
gravity 

Resistivity 

Temp, 
coefficient 
for  1°  C. 

Gram 
equivalents 
per  liter 

5 

1.0242 

2.55 

0.00159 

1.408 

10 

1.0490 

1.59 

0.00157 

2.884 

15 

1.0744 

1.35 

0.00156 

4.431 

20 

1.1001 

1.32 

0.00155 

6.050 

25 

1  .  1262 

1.39 

0.00154 

7.741 

30 

1  .  1524 

1.52 

0.00153 

9.506 

35 

1.1775 

1.70 

0.00152 

11.33 

40 

1.2007 

1.95 

13.22 

KOH  at  15° 

4.2 

1.0382 

6.85 

0.00188 

0.619 

8.4 

1.0777 

3.69 

0.00187 

1.580 

12.6 

1.1177 

2.67 

0.00189 

2.515 

16.8 

1.1588 

2.20 

0.00194 

3.477 

21.0 

1.2088 

1.97 

0  .  00200 

4.534 

25.2 

1.2439 

1.86 

0.00210 

5.599 

29.4 

1.2908 

1.85 

0.00222 

6.778 

33.6 

1.3332 

1.92 

0.00237 

8.001 

37.8 

1.3803 

2.10 

0.00258 

9.319 

42.0 

1.4298 

2.39 

0.00284 

10.730 

KC1  at  18° 

2.4 

1.0135 

29.10 

0.00219 

0.33 

5.0 

1.0308 

14.63 

0.00202 

0.693 

8.0 

1.0519 

8.20 

0.00200 

1.13 

10.0 

1  .  0638 

7.42 

0.00189 

1.43 

15.0 

1.0978 

4.99 

0.00180 

2.26 

19.3 

1  .  1308 

3.83 

0.00171 

2.93 

20.0 

1.1335 

3.77 

0.00169 

3.05 

25.0 

1  .  1408 

3.59 

0.00167 

3.83 

KI  at  18° 

5.0 

1  .  0363 

29.76 

0  .  00206 

0.312 

10.0 

1.0762 

14.81 

0.00201 

0.650 

20.0 

.1679 

6.94 

0.00158 

1.410 

30.0 

.2730 

4.38 

0.00167 

2.307 

40.0 

.3966 

3.18 

0.00152 

3.374 

50.0 

.5450 

2.57 

0.00144 

4.666 

55.0 

.6300 

2.38 

0.00141 

5.418 

KCN  at  15° 

3.25 

1.0154 

19.10 

0  .  00208 

0.508 

,      6.5 

1.0316 

9.80 

0.00194 

1.031 

APPENDIX 


141 


TABLE  6. — RESISTIVITY  OF  ELECTROLYTES. — Continued.     KOHL- 

RAUSCH    &    HOLBORN 


G  .  Salt 
per  100  g. 
of  sol. 

Specific 
gravity 

Resistivity 

Temp, 
coefficient 

Gram 
equivalent 
per    liter 

AgNO3atl8° 

5 

1.0422 

39.47 

0.00219 

0.307 

10 

1.0893 

21.20 

0.00218 

0.642 

15 

1  .  1404 

14.78 

0.00216 

1.009 

20 

1  .  1958 

11.57 

0.00213 

1.410 

25 

1.2555 

9.53 

0.00211 

1.851 

30 

1.3213 

8.14 

0.00210 

2.338 

35 

1.3945 

7.17 

0.00208 

2.879 

40 

1.4773 

6.45 

0.00206 

3.485 

.  45 

1.5705 

5.88 

0.00205 

4.168 

50 

1  .  6745 

5.44 

0.00206 

4.940 

55 

1.7895 

5.09 

0.00207 

5.800 

60 

1.9158 

4.80 

0.00210 

6.780 

KC2H3O2atl8° 

5 

1  .  0228 

29.03 

0.00224 

0.522 

10 

1.0466 

16.10 

0.00220 

1.069 

20 

1.0960 

9.62 

0.00223 

2.239 

30 

1  .  1484 

8.01 

0.00232 

3.519 

40 

1.2028 

7.97 

0.00251 

4.910 

50 

1.2598 

8.97 

0.00277 

6.430 

60 

1.3152 

11.94 

0.00325 

8.060 

70 

1.3714 

21.06 

0.00411 

9.810 

CuSO4  at  18° 

2.5 

1.0246 

92.4 

0.00214 

0.322 

5 

1.0513 

53.2 

0.00217 

0.661 

10 

1  .  1073 

31.4 

0.00219 

1.393 

15 

1  .  1675 

23.8 

0.00232 

2.202 

17.5 

1.2003 

21.9 

0.00237 

2.642 

142 


APPENDIX 


TABLE  7. — DEPOSITION     BY     IMMERSION. 
DEPOSITION,  PAGE  7 


GORE'S      ELECTRO- 


Solution 

Deposits  on 

Does  not  deposit  on 

SbCls                           Bi,  brass,  Germ.  Ag, 

Sb,  Cu,  Fe,  Ni,  Au, 

Pb,  Sn,  Zn. 

Pt,  Ag. 

Bids 

Fe,  Pb,  Sn,  Zn. 

Sb,  Bi,  brass,  Cu,  Au, 

Pt,  Ag. 

CuSO4,  Cu(NO3)2 

Fe,  Pb,  Sn,  Zn. 

Sb,  ,Bi,  Cu,  Au,  Ni, 

Pt," 

CuCl2 

Bi,  Fe,  Pb,  Sn,  Zn. 

Sb,  Cu,  Au,  Ni,  Pt, 

Ag. 

Ammoniacal, 

Zn. 

Sb,  Bi,  Cu,  Fe,  Au, 

CuCl2. 

Pb,  Ni,  Pt,  Ag. 

HgNOs 

As,  Bi,  Cd,  Cu,  Sb, 

Fe,  brass,  Pb,  Zn. 

AgN03 

Pb,  Sn,  Cd,  Zn,  Cu, 

Ag,  Au,  Pt. 

Bi,  Sb,  Fe,  Ni. 

Alcoholic  AgNO3 

As,  Sb,  Bi,  Zn,  Sn, 

Fe. 

Cu. 

AgCN.KCN 

Zn,    Pb,    Cu,   brass, 

Sb,  Bi,  Sn,  Fe,  Ni, 

Ger.  Ag. 

Ag,  Au,  Pt. 

Au(CN)3KCN 

Zn,   Cu,  brass,   Ger. 

Sb,  Bi,  Sn,  Pb,  Fe, 

Ag. 

Ni,  Ag,  Au,  Pt. 

APPENDIX 


143 


TABLE  8. — POTENTIALS 


OF   METALS  IN 

NEUMANN 


THEIR   NORMAL  SALTS. 


Sulphate 

Chloride 

Nitrate 

Acetate 

Magnesium 
Aluminum 
Manganese 
Zinc 
Cadmium 

+  1.239 
+  1.040 
+0.815 
+0.524 
+0.162 

+  1.231 
+  1.015 

+0.824 
+0.503 
+0.174 

+  1.060 
+0.775 
+0.560 
+0.473 
+0.122 

+  1.240 

+0.522 

Iron 
Cobalt 
Nickel 

+0.093 
-0.019 
-0.022 

+0.087 
-0.015 
-0.020 

-0.078 
-0.060 

-0.004 

Tin 

-0.085 

Lead 
Hydrogen 

-0.238 

-0.095 
-0.249 

-0.115 

-0.079 
-0.150 

Bismuth 
Antimony 

-0.490 

-0.315 
-0.376 

-0.500 

Arsenic 

-0  550 

Copper 
Mercury 

-0.515 
-0  980 

-0.615 
-1  028 

-0.580 

Silver 
Palladium 

-0.974 


-  1  .  066 

-1.055 

-0.991 

Platinum  "  ,. 

-1.140 

Gold 

-1.356 

TABLE  9. — ELECTRODE  POTENTIALS — CALCULATED  BY  THOMSON'S 
RULE.     WILSMORE 


K 

(+3.20) 

(+2.92) 

Pb 

+0.148 

-0.129 

Na 

(+2.82) 

(+2.54) 

H 

0.0 

-0.277 

Ba 

(+2.82) 

(+2.54) 

Cu 

-0.329 

-0.606 

Sr 

(+2.77) 

(+2.49) 

As 

<  -0.293 

<  -0.570 

Ca 

(+2,56) 

(+2.28) 

Bi 

<  -0.391 

<  -0.668 

Mg 

(+2.54) 

(+2.26) 

Sb 

<  -0.466 

<  -0.743 

Mg 

+  1.491? 

+  1.214? 

Hg 

-0.750 

-1.027 

Al 

+  1.276? 

+0.999? 

Ag 

-0.771 

-1.048 

Mn 

+  1.075  !      +0.798 

Pd 

<  -0.789 

<  -1.066 

Zn              +0.770 

+0.493 

Pt 

<  -0.863 

<-1.140 

Cd 

+0.420 

+0  .  143 

Au 

<  -1.079 

<  -1.356 

Fe 

+0.340 

+0.063 

Cl 

-1.417 

-1.694 

Co 

+0.232 

-0.045 

Br 

-0.993 

-1.270 

Ni 

+0.228 

-0.049 

I 

-0.520 

-0.797 

Sn 

<  +0.192 

<  -0.085 

O 

-1.119? 

-1.396? 

<  indicates  less  than  normal  ionic  concentration. 


144 


APPENDIX 


TABLE  10. — APPROXIMATE    POTENTIALS    IN    VARIOUS    ELECTRO- 
LYTES 13 


N.NaCl 

N.KCN 

^K2Cr20 

N.K2S 

N.NH4- 
CNS 

N.(NH4)2- 
C2O4 

N.KOH 

Mg  +  1.14|Al      +1.06  Mg   +.98Mg   +.80 

Mg  +  1.19 

Zn     +.65JA1   +1.06 

Mn        .60 

Zn           .90 

Mn       .35Cu      +.41 

Mn       .65 

Mn       .62Mg        .90 

Zn          .53 

Mg          .86 

Zn          .29  Mn          .32 

Zn         .60Cd         .30  Zn         .90 

Al           .30 

Cu           .81 

Al         .14Cd          .3lCd        .40 

Ai          .24Sn         .63 

Pb     +.04 

Mn          .65 

Cd    +.08 

Ag           .20iFe         .15 

Pb         .07Mn        .42 

Fe      -  .  04 

Cd           .64;Pb     -.08 

Zn           .19Sn         .02 

Sn     +.06Cd         .36 

Sn      -.07 

Monel    .38Sn     -.15 

Al            .15 

Cu    +  .01N1     -.05:Pb         .28 

Cr      -.17 

Ni           .  36  Fe     -  .  30 

Sn           .11 

Pb         .00 

Nichrome  Fe         .27 

-.08 

Bi      -  .  24 

Au           .34Sb     -.31 

Cr           .07 

Ni     -.07 

Cr     -.19 

Bi          .04 

Sb      -  .  29 

Sn           .30Cu     -.40 

Au           .05 

Si      -.24 

W-    -.21 

Mo   +.01 

Cu     -.31 

Nichrome     Mo    —.40 

Ni           .05 

Ag     -.25 

Fe     -.26 

Si      -.04 

.30 

Mo    -.35 

Ag           .28|Si      -.45 

Monel    .04 

Sb     -  .  28 

Cu    -.30 

Ni     -  .  06 

Ni      -.37 

Pb           .15 

Bi      -.51 

Te       +  .  03 

Bi      -.34 

Bi      -.31 

Cr     -.10 

W      -.38 

Sb           .13 

Ag     -.62 

Pt            .03 

Mo    -  .  37 

Monel 

Cu    -.14 

-.34 

Ag     -  .  50 

Hg          .12 

Cr     -.62 

Bi            .03 

Pt     -.42 

Mo    -.38 

C       -.25 

Te      -.62 

Cr           .11 

Hg    -.62 

Nichrome 

Cr     -.56 

Ag     -.48 

Ag     -.33 

+  .03 

Au     -  .  62 

Si        +.07 

W      -.62 

Pb       +.03 

C       -.60 

C       -.48 

Au    -.33 

Pt      -  .  63 

Bi       -.08 

Nichrome 

f*a 

W        +.02 

Te     -.53 

Pt     -  .  37 

C       -  .  63 

W        -.08 

—  .  DO 

Monel  .67 

C         -.01 

Au    -.53 

Pt       -.08 

Ni     -  .  74 

Si        -.01 

Pt     -  .  54 

Fe       -  .  12 

C       -.81 

Fe       -  .  07 

Te       -.13 

Pt     -.93 

Mo     -.11 

Mo      -  .  14 

Te     -.93 

PbO2  -.60 

C         -.36 

Au    -  .93 

PbO2  -.68 

13  The  values  given  in  this  table  are  the  average  results  obtained 
by  students,  and  are  probably  seriously  in  error  in  many  instances. 
Although  the  order  of  the  elements  is  more  reliable  than  the  nu- 
merical values,  even  this  cannot  be  absolutely  relied  upon.  The 
author  will  appreciate  corrections  and  additions  to  the  data  set 
forth  in  the  table. 


APPENDIX 


145 


TABLE  11. — DECOMPOSITION  VOLTAGE.     LE  BLANC 


H2S04 

1.67  volt 

SrCl2 

2.01 

HNO3 

1.69 

BaCl2 

1.95 

H3P04 

1.70 

ZnSO4 

2.35 

HC1 

1.31 

ZnBr 

1.80 

NaOH 

1.67 

NiSO4 

2.09 

KOH 

1.69 

NiCl2 

1.84 

NH4OH 

1.74 

AgN03 

0.70 

Na2SO4 

2.21 

CdSO4 

2.03 

NaNO3 

2.15 

CoSO4 

1.92 

NaCl 

1.98 

HgCl2 

1.30 

NaBr 

1.58 

Fe2(SO4)3 

1.62 

Nal 

1.12 

FeSO4 

2.02 

NaC2H3O2 

2.10 

AuCl3 

0.39 

K2SO4 

2.20 

FeCl2 

2.16 

KNO3 

2.17 

SnCl2 

1.76 

KC1 

1.96 

MnSO4 

2.60 

(NH4)2S04 

2.11 

MnCl2 

2.77 

CaCl2 

1.89 

CuCl2 

1.36 

INDEX 


Addition  agents  84,  89 
Alloys,  electrodeposition  of,  65 

97,  101,  113-115 
Aluminum  electrodes,  29 

corrosion  of,  70,  112 
electroplating  of,  110,  111 
rectifier,  29 

Ammonium  persulphate,  126 
Atomic  weight,  table  of,  138 

Barium  perchlorate,  131 
Brass  plating,  60,  66,  97-101, 
107,  108,  113 

Calomel  electrode,  36,  39 
Chlorates,   preparation  of,  127, 

128 

Coloring  metals,  118,  119 
Copper     plating,    95-97,    104, 

110,  111,  113 

Corrosion  of  metals,  60-62 
Coulombmeter,  30,  31,  111 
Critical  current  density,  107 
Current  efficiency,  30,  111,  112 

Decomposition  voltage,  see 
"E.M.F.  of  decom- 
position. 

Deposition  by  immersion,  103 
of  metals,  see  "Electrode- 
position." 

Discharge  potential,  56-57 

Electrical     instruments,     prin- 
ciples of  construction, 
7 
protection  from  injury,  8 


Electric  cleaning  of  metals,  60, 

91 
Electrochemical       equivalents, 

138. 
Electrode,    aluminum,    29,    70, 

112 

intermediate,  69 
normal  calomel,  36,  39 
Electro-deposition  of  alloys,  97- 

101 

of  brass,  66,  97-101,  113 
of  cadmium,  116 
of  copper,  68,  95-97,  104, 

110,  111,  113 
of  lead,  12,  113 
of  iron,  116 

of  nickel,  68,   93-95,   104, 
105,     108,     109,     111, 
112 
of  silver,  97,  101,  102,  105, 

115 

principles  of,  84-89 
Electrolytic  analysis,  66 

oxidation,     121,  122,     124- 

132,  133,  135,  136 
reduction,  11,  33,  121-123 
separation  of  metals,  63 
Electromotive  force,  33,  34 
of  decomposition,  43 
of  copper  sulphate,  49 
of  hydrochloric  acid,  48, 

50 

of  nitric  acid,  48,  50 
of  sodium  chloride,  43 
of  sulphuric  acid,  47,  48, 

50 

of  zinc  bromide,  46,  50 
147 


148 


INDEX 


Electromotive   force,    effect  of 
electrodes     of     unequal 
size  upon,  49-55 
table  of  numerical  values, 
145 

Faraday's  law,  29 
Fig.    1 — Resistivity  of  electro- 
lytes, 14 

2 — Resistivity,     using    a 
low-scale  voltmeter,  17 
3 — Resistivity     using     a 
double-scale    volt- 
meter, 18 

4 — Resistivity,  of  wire,  19 
5 — Resistivity,  with  volt- 
meter only,  21 
6 — Resistivity,   by    volt- 
meter  and    resistance 
box,  22 

7 — Cell    for    temperature 
coefficient     of     resist- 
ance, 23 
8 — Resistance  of  a  fused 

electrolyte,  25 
9 — Potentiometer         and 

connections,  38 
10 — Polarization  in  a  vol- 
taic cell,  42 

11 — E.M.F.  of  decomposi- 
tion, 43 

12 — E.M.F.  of  decomposi- 
tion and  polarization, 
45 

13 — E.M.F.  of  decomposi- 
tion  and   polarization 
by  the  potentiometer, 
51 
14 — Curves  of  polarization 

at  electrodes,  53 
15 — Curves  for  electrodes 
of  unequal  size,  54 


Fig.  16 — Apparatus  for  electro- 
lytic analysis,  67 

Galvanite  plating,  105 
Galvanometer  key,  special,  52 
Galvanoplasty,  119 
Gas     volumes,     reduction     to 

standard      conditions, 

32 

Hypochlorite  of  sodium,  135 
Instruction  for  students,  5 

Laboratory  equipment,  2 
Lead  nitrate,  electrolysis  of,  12 

Mercury  cathode,  12 

refining  of,  134 
Metallochromes,  118 

Nickel  plating,  93-95,  104,  105, 
108,  109,  111,  112 

Ohm's  law,  14 
Overvoltage,  58 

of  hydrogen,  58,  121 
of 'oxygen,  122 
Oxidation,     121,  122,     129-132, 

133,  135,  136 
Oxidizing  metals,  baths  for,  80- 

82 
experiments,  118-119 

Passive  state  of  metals,  58-60 
Perchlorates,  130,  131 
Persulphates,  estimation  of,  126 

preparation  of,   125-127 
Plating  baths,  composition  of, 

72-79 

metal  content  and  cur- 
rent density,  83 


INDEX 


149 


Plating  by  contact,  104 

by  immersion,  103 

general  principles,  84-89 
Polarization,  15 

effect  of  size  of  electrodes 
on,  49,  55 

measurement  of,  17,  44-48 
Polishing  metals,  92 
Potassium,  137 

bromate,  128 

bromide,  11,  128 

chlorate,  127 

chloride,     electrolysis     of, 
63,  127 

chromate,  133 

ferrocyanide,  136 

iodate,  129 

perchlorate,  130 

permanganate,  133 

persulphate,  125 
Potential,  33-41 

of  electrodes,  measurement 

of,  39-41 
Potentiometer,  37 

Quicking,  101 

Reduction,  electrolytic    11,  33, 

121-123 
Resistance,     measurement     of 

13-22 

of  a  fused  electrolyte,  25 
of    glass,    effect    of    heat 

upon,  26 

of  primary  cells,  28 
of  storage  cells,  28 
of  a  voltmeter,  20 
temperature  coefficient  of, 

19,  23 
Resistivity,  13 

effect  of  the  solvent  upon, 
24 


Resistivity  of  copper  sulphate, 

15,  17,  19 
of  metals,  18,  139 
of  plating  baths,  19 

Scheele's  green,  134 
Separation  of  metals,    electro- 
lytic, 63 
Single  potentials,   34-41,     143, 

144" 
Silver  plating,  97,  101,  102,  105, 

115 

striking  baths,  102 
Sodium     chloride,     electrolysis 

of,  11,  12,  16,  135 
hydrate,  electrolysis  of,  31 
hypochlorite,     preparation 

of,  135 
nitrate,  electrolysis  of,  11, 

33 

sulphate,  electrolysis  of,  63 
Specific  resistance,  see  "  Resis- 
tivity." 
Spotting  out  of  plated  articles, 

97 
Switch-key,  special,  52 

Table  of  atomic  weights,  138 

of  deomposition  voltage, 
145 

of  deposition  by  immersion, 
142 

of  electrochemical  equiva- 
lents, 139 

of  metal  content  of  plating 
baths,  83 

of    overvoltage    of   hydro- 
gen, 121 
of  oxygen,  122 

of  plating  baths  metal  con- 
tent vs.  current  den- 
sity, 83 


150  INDEX 

Table  of  potentials  calculated  by  Table  resistivity  of  metals,  139 

Thomson's  rule,   143  Tin  plating,  104 
in  different  electrolytes, 

144  Voltite  plating,  105 
in  normal  solutions,  142 

resistivity   of   electrolytes,  Water,  electrolysis  of,  31 

139  White  lead,  preparation  of,  12 


THIS  BOOK  IS  DUE  ON  THE  LAST  DATE 
STAMPED  BELOW 


AN  INITIAL  FINE  OF  25  CENTS 

WILL  BE  ASSESSED  FOR  FAILURE  TO  RETURN 
THIS  BOOK  ON  THE  DATE  DUE.  THE  PENALTY 
WILL  INCREASE  TO  SO  CENTS  ON  THE  FOURTH 
DAY  AND  TO  $1.OO  ON  THE  SEVENTH  DAY 
OVERDUE. 


MAR   10  1937 

ccp  7   J937  i 

Wm  \ 

*P*    12H39 

»% 

to-  ^ 

Y 

LD  21-100m-8,'34 

te 


lA/3 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 


